UNIVERSITY  OF  ILLINOIS 

CHEMISTRY  DEPARTMENT 


ARTHUR  WILLIAM  PALMER 
MEMORIAL  LIBRARY 
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A TEXT-BOOK 

OF 

INORGANIC  CHEMISTRY 


RICHTER 


Richter’S  Chemistries 

AUTHORIZED  TRANSLATIONS 


BY  EDGAR  F.  SMITH,  M.A.,  Rm.D.,  Sc.D. 

PROFESSOR  OF  CHEMISTRY  IN  THE  UNIVERSITY  OF  PENNSYLVANIA;  MEMIIER  OF  THE  CHEMICAL 
SOCIETY  OF  BERLIN,  AMERICAN  PHILOSOPHICAL  SOCIETY,  ETC.,  ETC. 


Professor  Richter’s  methods  of  arrangement  and  teaching  have  proved  their 
superiority  by  the  iarge  sale  of  his  books  throughout  Europe  and  America,  trans- 
lations having  been  made  in  Russia,  Holland,  and  Italy.  The  success  attending 
tlieir  publication  in  this  country  could  only  have  been  attained  for  good  books  that 
have  been  found  useful,  practical,  and  tlioroughly  up  to  the  times. 

The  Chemistry  of  the  Carbon  Compounds, 
or  Organic  Chemistry 

Third  American,  translated  from  the  Eighth  German  Edition  by 
Prof.  R.  Anschutz,  University  of  Bonn.  Thoroughly  Revised, 
Enlarged,  and  in  many  parts  Rewritten.  Illustrated.  Two  Volumes. 

Vol.  I.  Aliphatic  Series.  625  pages.  . Cloth,  net,  1^3.00 

Vol.  II.  Carbocyclic  and  Heterocyclic  Series.  671  pages. 

Cloth,  net,  $3.00 


INORGANIC  CHEMISTRY 

Fifth  American  from  the  Tenth  German  Edition  by  PROF.  H. 
Klinger,  University  of  Konigsberg.  Thoroughly  Revised  and  in 
many  parts  Rewritten.  With  many  Illustrations  and  Colored  Plate 
of  Spectra.  Cloth,  net,  J^i.75 


***  Special  Catalogues  of  Books  on  Chemistry,  Hygiene,  Pharmacy,  Medicine, 
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VICTOR  VON  RICHTER’S 


TEXT-BOOK 

OF 

INORGANIC  CHEMISTRY 

EDITED  BY 

PROF.  H.  KLINGER 

UNIVERSITY  OF  KOENIGSBERG 


AUTHORIZED  TRANSLATION 

BY 

EDGAR  F.  SMITH 


PROFESSOR  OF  CHEMISTRY  IN  THE  UNIVERSITY  OF  PENNSYLVANIA,  PHILADELPHIA 


(Assisted  by  WALTER  T.  TAGGART) 

INSTRUCTOR  IN  CHEMISTRY 


jfittb  Hmerican  trom  tbe  Uentb  (Berman  lEbition 

CAREFULLY  REVISED  AND  CORRECTED 


WITH  SIXTY-EIGHT  ILLUSTRATIONS  ON  WOOD 

AND 

COLORED  LITHOGRAPHIC  PLATE  OF  SPECTRA 


PHILADELPHIA 

P.  BLAKISTON’S  SON  & CO. 

1012  WALNUT  STREET 

1900 


Copyright,  1900,  by  P.  Blakiston’s  Son  & Co. 


WM.  F.  FELL  & CO., 
ELECTHOTVPEH8  AND  PHINTER8, 
1320-24  8ANS0M  8THEET, 
PHILADELPHIA. 


PREFACE 


TO  THE 


FIFTH  AMERICAN  EDITION. 


The  student  of  the  present  edition  will  discover  that  it  differs  very 
materially  from  all  preceding  editions.  This  is  largely  due  to  the  fact 
that  the  editor  has  endeavored  to  give  due  consideration  to  the  more 
recent,  well-established  discoveries  in  chemical  science  ; hence  additions 
will  be  found  relating  to  the  general  properties  and  the  measurement  of 
gases,  to  the  atmosphere  and  the  interesting  constituents  lately  observed 
in  it,  to  the  theory  of  dilute  solutions  and  electrolytic  dissociation,  to  the 
electrolysis  of  salts,  to  alloys,  etc.  Thus  revised,  it  is  hoped  that  the 
book  will  continue  to  occupy  the  position  it  has  so  long  held  among 
works  devoted  to  the  inorganic  portion  of  chemical  science. 

The  translator  would  take  this  opportunity  to  acknowledge  his  great 
indebtedness  and  to  return  his  sincere  thanks  to  Mr.  Walter  T.  Taggart, 
upon  whom  devolved  the  task  of  arranging  the  crude  manuscript  for  the 
press  and  the  revision  of  the  proof-sheets. 

The  John  Harrison 
Laboratory  of  Che77iistry. 


V 


PREFACE 


TO  THE 

FIRST  AMERICAN  EDITION. 


The  success  of  Prof,  von  Richter’s  work  abroad  would  indicate  its 
possession  of  more  than  ordinary  merit.  This  we  believe  true,  inasmuch 
as,  in  presenting  his  subject  to  the  student,  the  author  has  made  it  a 
point  to  bring  out  prominently  the  relations  existing  between  fact  and 
theory.  These,  as  well  known,  are,  in  most  text-books  upon  inorganic 
chemistry,  considered  apart,  as  if  having  little  in  common.  The  results 
attained  by  the  latter  method  are  generally  unsatisfactory.  The  first 
course — that  adopted  by  our  author — to  most  minds  would  be  the  more 
rational.  To  have  experiments  accurately  described  and  carefully  per- 
formed, with  a view  of  drawing  conclusions  from  the  same  and  proving 
the  intimate  connection  between  their  results  and  the  theories  based  upon 
them,  is  obviously  preferable  to  their  separate  study,  especially  when  they 
are  treated  in  widely  removed  sections  or  chapters  of  the  same  book. 
Judging  from  the  great  demand  for  von  Richter’s  work,  occasioning  the 
rapid  appearance  of  three  editions,  the  common  verdict  would  seem  to 
be  unanimously  in  favor  of  its  inductive  methods. 

In  the  third  edition,  of  which  the  present  is  a translation,  the  Periodic 
System  of  the  Elements,  as  announced  by  Mendelejeff  and  Lothar 
Meyer,  is  somewhat  different,  in  the  manner  of  development  and  pre- 
sentation, from  that  appearing  in  the  previous  editions.  This  was  done 
to  give  more  prominence  to  and  make  more  general  the  interesting  rela- 
tions disclosed  by  it.  Persons  examining  this  system  carefully  will  be 
surprised  to  discover  what  a valuable  aid  it  really  has  been,  and  is  yet,  in 
chemical  studies.  Through  it  we  are  continually  arriving  at  new  rela- 
tions and  facts,  so  that  we  cannot  well  hesitate  any  longer  in  adopting  it 
into  works  of  this  character.  It  is,  indeed,  made  the  basis  of  the 

vii 


viii  PREFACE  TO  THE  FIRST  AMERICAN  EDITION. 

present  volume.  In  accordance  with  it,  some  change  in  the  treatment 
of  the  metals,  ordinarily  arbitrarily  considered,  has  been  made. 

A new  feature  of  the  work,  and  one  essentially  enlarging  it,  is  the 
introduction  of  the  thermo-chemical  phenomena,  briefly  presented  in  the 
individual  groups  of  the  elements  and  in  separate  chapters,  together  with 
the  chemical  affinity  relations  and  the  law  of  ])eriodicity.  “Hereby 
more  importance  is  attributed  to  the  ])rinciple  of  the  greatest  heat 
development  than  at  present  apjiears  to  belong  to  it,  because  it  was 
desired,  from  didactic  considerations,  by  the  explanation  of  the  few 
anomalies,  to  afford  the  student  the  incentive  and  o[)])ortunity  of  deduc- 
tively obtaining  the  majority  of  facts  from  the  thermal  numbers,  on  the 
basis  of  a simple  principle,  do  facilitate  matters,  there  is  appended  to 
the  volume  a table  containing  the  heat  of  formation  of  the  most  im- 
portant.compounds  of  the  metals.” 

Trusting  that  the  teachings  of  this  work  will  receive  a hearty  welcome 
in  this  country,  and  that  they  will  meet  a want  felt  and  often  expressed 
by  students  and  teachers,  we  submit  the  following  translation  of  the 


same. 


TABLE  OF  CONTENTS 


INTRODUCTION. 

Physics  and  Chemistry,  17.  Physical  and  Chemical  Phenomena,  1 8,  19.  Chemical 
Elements,  19.  Principle  of  Indestructibility  of  Matter,  20.  Principle  of  Conserva- 
tion of  Energy,  21.  Forms  and  Equivalents  of  Energy,  22.  Chemical  Energy,  23. 
Constitution  of  Matter ; Atom  and  Molecule,  24.  Chemical  Symbols  and  For- 
mulas, 25.  Atomic  Weights,  26.  Table  of  Atomic  Weights,  26.  Chemical  Equa- 
tions, 27.  Conditions  of  Chemical  Action,  28.  T hermo-chemical  Phenomena,  28. 
Crystallography,  31. 


HYDROGEN  AND  THE  NON-METALS. 

Classification  of  the  Elements,  39. 

Hydrogen,  40.  Purifying  and  Drying  of  Gases,  42.  Apparatus  for  the  Generation 
and  Collection  of  Gases,  42.  Physical  Properties  of  Hydrogen,  43.  Chemical 
Properties  of  Hydrogen,  45.  Condensation  of  Gases,  47.  Critical  Condition,  47. 
Group  of  Halogens,  49. 

Chlorine,  49,  Bromine,  53.  Iodine,  54.  Fluorine,  56.  General  Characteristics 
of  the  Halogens,  56.  Compounds  of  the  Halogens  with  Hydrogen,  57.  Hydrogen 
Chloride,  57.  Acids  ; Bases  ; Salts,  60.  Hydrogen  Bromide,  61.  Hydrogen  Iodide, 
63.  Hydrogen  P'luoride,  64.  General  Characteristics  of  the  Hydrogen-halogen 
Compounds,  65.  Thermo-chemical  Deportment  of  the  Halogens,  66.  Compounds 
of  the  Halogens  with  one  another,  68. 

Weight  Proportions  in  the  Union  of  the  Elements;  Stoechiometric  Laws;  Atomic 
Hypothesis;  Choice  of  Atomic  Weights,  69.  General  Properties  of  Gases;  Atomic- 
molecular  Theory,  73.  Avogadro’s  Law,  75.  Status  nascens,  77.  Principles  of 
the  Atomic-molecular  Theory,  79.  Determination  of  the  Atomic  Weights,  80, 
Oxygen  Group,  80. 

Oxygen,  80.  Oxyhydrogen,  83.  Oxidation  and  Reduction,  84.  Ozone,  84. 
Isomerism  and  Allotropy,  87.  Compounds  of  Oxygen  with  Hydrogen,  88.  Water, 
88.  Natural  Waters  ; Chemical  Properties  of  Water,  91.  Electrolysis  of  Water  ; 
Thermo-chemical  Deportment,  92.  Dissociation,  93.  Kinetic  Theory  of  Gases,  94. 
Quantitative  Composition  of  Water,  95.  Molecular  Formula  of  Water  ; Atomic 
Weight  of  Hydrogen  and  of  Oxygen,  97.  Hydrogen  Peroxide,  99.  Catalysis; 
Thermo-chemical  Dejiortment,  103. 

Sulphur,  104. 

Molecules  of  the  Elements,  106.  Hydrogen  Sulphide,  107.  Hydrogen  Persulphide, 
1 10.  Compounds  of  Suljihur  with  the  Halogens,  no.  Selenium,  112.  Tellurium, 
1 1 3.  Summary  of  the  Elements  of  the  Oxygen  Group,  113.  Thermo-chemical 
Deportment,  114. 


IX 


X 


TAI’.LF,  OF  CONTENTS. 


Nitrogen  Group,  114. 

Nitrogen,  115.  Alinospliere,  n6.  luKlioinetry,  120.  Measuring  Gases,  121. 
Diffusion  of  (jases,  123.  Gases  Recently  Discovered  in  the  Atmosphere,  123. 
Argon,  124.  1 leliuin,  125.  Goinpouiufs  of  Nitrogen  with  Hydrogen,  125.  Atn- 

monia,  126,  Amtnoniuin  Salts,  129.  Atoinie  Weight  of  Nitrogen,  129.  Ilydrox- 
ylainine,  130,  Dianiide,  131.  Ilydrazoie  Aeid,  132.  Compounds  of  Nitrogen 
with  tlie  Halogens,  133.  IMiosphorus,  135.  Compounds  of  J'hosphorus  with 
Hydrogen,  138.  I'hosphoniuin  Salts,  140.  Comi)ounds  of  IMiosphorus  with  the 
Halogens,  141.  Arsenie,  143.  Arsine,  144.  Compounds  of  Arsenie  with  the 
Halogens,  145.  Antimony,  146,  Stihine,  147.  Comj)ounds  of  Antimony  with 
the  Halogens,  147.  Tabulation  of  the  Elements  of  the  Nitrogen  Grouji,  148. 

Carbon  Group,  149. 

Carbon,  149.  Carbon  Compounds  of  Hydrogen,  15 1.  Methane,  152,  Atomic 
Weight  of  Carbon,  152.  Ethane;  Ethylene,  153.  Acetylene,  154.  Nature  of 
Elame,  155.  Comixninds  of  Carbon  with  the  Halogens,  159.  Silicon,  iCo. 
Hydrogen  Silicide,  161.  Compounds  of  Silicon  with  the  Halogens,  161,  162. 
Hydrogen  Silico-fluoride,  163.  Silicon  Carbide  (Carborundum),  163. 

Valence  of  the  Elements;  Chemical  Structure  of  the  Molecules,  165. 

Oxygen  Compounds  of  the  Metalloids,  172. 

Oxygen  Compounds  of  the  Halogens,  173. 

Oxygen  Compounds  of  Chlorine,  172.  Ilypochlorous  Oxide,  174.  Hypochlorous 
Acid,  174.  Chlorine  Trioxide  ; Chlorous  Acid  ; Chlorine  Tetroxide,  176.  Chloric 
Acid,  177.  rerchloric  Acid,  178.  Oxygen  Compounds  of  Eromine,  179.  Oxygen 
Compounds  of  Iodine,  180.  Hydrate*  of  the  Acids,  181. 

Oxygen  Compounds  of  the  Elements  of  the  Sulphur  Group,  182. 

Oxygen  Compounds  of  Sulphur,  183.  Sulphur  Dioxide,  183.  Sulphurous  Acid, 
185.  Hydrosulphurous  Acid,  186.  Sulphur  Sesquioxide,  187.  Sulphur  Trioxide, 

187.  Thermo-chemical  De])ortment,  187.  .Sulphur  Heptoxide  ; Persulphuric  Acid, 

188.  Sulphuric  Acid,  189  Disulphuric  Acid,  193.  Sulphuric  Acid  Chloranhydrides, 
195.  Chlorsulphonic  Acid;  Sulphuryl  Chloride,  195.  Amido-derivatives,  196. 
Thiosulphuric  Acid  ; Polythionic  Acids,  197.  Oxygen  Derivatives  of  Selenium  and 
Tellurium,  199. 

Oxygen  Compounds  of  the  Elements  of  the  Nitrogen  Group,  200. 

Oxygen  Derivatives  of  Nitrogen,  201.  Nitric  Acid;  Nitrogen  Pentoxide  ; Nitryl 
Chloride;  Nitrosyl  Chloride;  Nitramide,  204.  Nitrogen  Trioxide  ; Nitrous  Acid, 
205.  Nitrogen  Tetroxide,  206.  Nitrosyl-sulphuric  Acid,  207.  Nitric  Oxide,  209. 
Nitrous  Oxide,  21 1.  Hyponitrous  Acid,  212.  Compounds  of  Nitrogen  and  Sul- 
phur, 213.  Oxygen  Compounds  of  Phosphorus,  213.  Hypophosphorous  Acid,  214. 
Phosphorous  Acid,  215.  Phosphoric  Acid  ; Pyrophosphoric  Acid  ; Hypophosphoric 
Acid,  216.  Metaphosphoric  Acid  ; Phosphorus  Pentoxide,  217.  Chloranhydrides 
of  the  Acids  of  Phosphorus,  218.  Phosphorus  Compounds  with  Sulphur,  219. 
Oxygen  Derivatives  of  Arsenic,  220.  Arsenic  Trioxide,  220.  Arsenic  Acid,  221. 
Compounds  of  Arsenic  with  Sulphur,  222.  Sulpho-salts,  223.  Oxygen  Derivatives 
of  Antimony,  223.  Antimony  Oxide  ; Antimonic  Acid,  224.  Antimony  Sulphides, 
225.  Vanadium  ; Niobium  ; Tantalum,  226. 

Oxygen  Compounds  of  the  Elements  of  the  Carbon  Group,  226. 

Oxygen  Compounds  of  Carbon,  227.  Carbon  Dioxide,  227.  Critical  Pressure,  228. 
Physiological  Importance  of  Carbon  Dioxide  ; Percarbonic  Acid,  231.  Carbon  Mon- 
oxide, 231.  Nickel  Carbonyl,  233.  Carbonyl  Chloride,  234.  Amido-derivatives  of 
Carbonic  Acid,  234.  Compounds  of  Carbon  with  Sulphur,  234.  Cyanogen  Com- 
pounds, 235.  'rhermo-chemistry  of  the  Carbon  Compounds,  236.  Oxygen  Com- 
pounds of  Silicon,  236.  Dialysis,  237.  Crystalloids  and  Colloids,  238.  Silicates, 
238.  M'itanium  ; Zirconium;  M'horium,  238.  lk)ron,  24I.  Boron  Hydride,  241. 
Honm  (Chloride,  241.  Boron  Eluoride,  242.  Boric  Acid,  242. 

The  Periodic  System  of  the  Elements,  243. 

I’criodicily  of  Chemical  Valence,  2|8.  Correction  of  Atomic  Weights,  250. 


TABLE  OF  CONTENTS. 


XI 


THE  METALS. 

Physical  Properties  of  the  Metals,  251.  Atomic  Volumes,  252.  Light  and  Heavy 
Metals,  252.  Melting  Points  of  the  Metals,  252.  Electric  Furnace,  253.  Specific 
Heat;  Atomic  Heat,  253.  Thermal  Atomic  Weights,  254.  Isomorphism,  255. 
Chemical  Properties  of  the  Metals,  255.  Alloys,  255.  Amalgams,  257.  Metallic 
Carbides,  257.  Halogen  Compounds,  257.  Oxides  and  Hydroxides,  258.  Per- 
oxides, 259.  Salts,  259.  Action  of  Metals  on  Salts  and  Acids,  261.  Electrolysis 
of  Salts,  262.  Faraday’s  Law,  263.  Solutions,  265.  Theory  of  Dilute  Solutions, 
267.  Theory  of  Electrolytic  Dissociation,  268.  Transposition  of  Salts,  270. 

Group  of  the  Alkali  Metals,  272. 

d'hermo-chemistry  of  the  Alkali  Metals,  273.  Potassium,  273.  Potassium  Hydride, 
274.  Dissociation,  274.  Potassium  Oxides,  275.  Potassium  Hydroxide,  275. 
Potassium  Chloride,  276.  Bromide,  276.  Iodide,  276.  Fluoride,  277.  Cyanide, 
277.  Potassium  Chlorate,  277.  Potassium  Hypochlorite,  278.  Potassium  Sulphate, 
279.  Potassium  Sulphite,  279.  Potassium  Persulphate,  279.  Potassium  Nitrate, 
279.  Gunpowder,  280.  Potassium  Nitrite,  280.  Potassium  Phosphate,  280.  Potas- 
sium Carbonate,  280.  Potassium  Silicate,  282.  Potassium  .Sulphides,  283.  Potas- 
sium Amide,  283.  Recognition  of  the  Potassium  Compounds,  283.  Rubidium  ; 
Caesium,  284.  Sodium,  284.  Sodium  Oxides,  285.  Sodium  Hydroxide,  286. 
Sodium  Chloride,  286.  Sodium  Bromide,  287.  Sodium  Iodide,  287.  Sodium 
Chlorate,  287.  Sodium  lodate,  287.  Sodium  Sulphate,  287.  Supersaturated 
Solutions,  287.  Sodium  Hyposulphite,  289.  Sodium  Carbonate,  289.  Sodium 
Nitrate  ; Sodium  Phosphates,  293.  Borax,  294.  Sodium  Silicate,  295.  Sodium 
Nitride,  295.  Recognition  of  Sodium  Compounds,  295.  Lithium,  295.  Ammo- 
nium Compounds,  296.  Ammonium  Chloride,  297.  Ammonium  Carbonate,  298. 
Ammonium  Phosphates,  298,  Ammonium  Nitride,  298.  Ammonium  Sulphide, 
299.  Ammonium  Hydrosulphide,  299.  Recognition  of  Ammonium  Compounds, 
299- 


METALS  OF  GROUP  II,  299. 

Group  of  the  Alkaline  Earths,  3C0. 

Calcium,  300.  Calcium  Oxide,  301.  Cement,  302.  Calcium  Chloride,  302.  Cal- 
cium Fluoride,  302.  Chloride  of  Lime,  303.  Calcium  Sulphate,  304.  Calcium 
Phosphates,  304.  Calcium  Carbonate,  305.  Glass,  306.  Calcium  Sulphides,  307. 
Calcium  Carbide,  307.  Strontium,  307.  Barium,  308.  Barium  Oxide,  308.  Barium 
Peroxide,  309.  Barium  Sulphate,  309.  Barium  Persulphate,  309.  Barium  Car- 
bonate, 310.  Recognition  of  the  Compounds  of  the  Alkaline  Earths,  310.  Diam- 
monium Compounds,  310.  Hydrazine  Hydrate,  31 1.  Diammonium  Chloride,  312. 
Diammonium  Nitride,  312.  Azides,  312. 

Magnesium  Group,  312. 

Magnesium,  313.  Magnesia,  314.  Magnesium  Chloride,  314.  Magnesium  Sul- 
phate, 315.  Magnesium  Phosphates,  315.  Magnesium  Carbonate,  316.  Mag- 
nesium Nitride,  316.  Recognition  of  Magnesium  Compounds,  317.  Beryllium, 
317.  Zinc,  318.  Zinc  Oxide,  318.  Zinc  Chloride,  319.  Zinc  Sulphate,  319. 
Zinc  Sulphide,  319.  Cadmium,  320.  Comparison  of  Zinc,  Cadmium,  and  Mercury, 
321.  Mercury,  322.  Amalgams,  323.  Mercurous  Compounds,  323.  Mercuric 
Compounds,  325.  Heat  of  Formation  of  the  Chlorides  of  Group  II,  327. 

Copper,  Silver,  Gold,  328.  General  Characteristics,  329.  Forms  of  Combination,  329. 
Copper,  330.  Metallurgy  of  Copper,  331.  Cuprous  Compounds,  332.  Cupric  Com- 
pounds, 334.  Alloys  of  Copj)er,  335.  Silver,  336.  Metallurgy,  337.  Silver  Oxides, 
338.  Molecular  Formulas,  339.  Silver  Chloride,  339.  Photography,  340.  Nitrate 
of  Silver,  341.  Silver  Nitride,  341,  Silver  Sulphide,  341.  Silvering,  342.  Gold, 
342.  Aurous  Compounds,  343.  Auric  Compounds,  344. 


Xll 


TABLK  OF  CONl'ENTS. 


METALS  OF  GROUP  III,  345. 

Group  of  Earth  Metals,  346. 

Aluminium,  346,  Aluminium  Chloride,  348.  Aluminium  Oxide,  350.  Aluminates, 
350.  Aluminium  Sulphate,  352.  Alum,  352.  Aluminium  Silicates,  353.  Porcelain, 
353.  Ultramarine,  354.  Rare  Earth  Metals,  354. 

Gallium  Group,  357. 

Gallium,  358.  Indium,  358.  Thallium,  359.  'I'hallous  Compounds,  360,  Thallic 
Compounds,  360. 

METALS  OF  GROUP  IV,  361. 

Germanium,  362.  Germanous  Compounds,  362.  Germanic  Compounds,  363. 

Tin,  363.  Stannous  Compounds,  364.  Stannic  Comi)ounds,  366.  Stannates,  367. 
Sulpho-stannates,  367.  Lead,  367.  Lead  Oxide,  369.  Plumbic  Acid,  370.  Lead 
Chloride,  370.  Lead  Nitrate,  371.  Lead  Carbonate,  371.  Lead  Sulphide,  371. 
Bismuth,  372.  Pismuthic  Acid,  373.  Pismuth  Nitrate,  373. 

Chromium  Group,  373. 

Chromium,  374.  Chromous  Compounds,  375.  Chromic  Compounds,  375.  Chro- 
mium Alum,  377.  Chromates,  377.  Potassium  Chromate,  379.  Chromyl  Chloride, 
380.  Molybdenum,  382.  Tungsten,  384.  Uranium,  384. 

Manganese,  386. 

Forms  of  Combination,  386.  Manganous  Compounds,  387.  Manganic  Compounds, 
388.  Manganese  Peroxide,  389.  The  Acids  of  Manganese,  390. 

METALS  OF  GROUP  VIII,  392. 

Iron  Group,  392. 

Iron,  393.  Cast  Iron,  394.  Wrought  Iron,  394.  Metallurgy  of  Iron,  395.  Fer- 
rous Compounds,  397.  Ferric  Compounds,  399.  Ferric  Acid  Compounds,  401. 
Cyanogen  Compounds,  401.  Metallic  Ions,  401.  Nickel,  404.  Cobalt,  406. 
Platinum  Metals,  408. 

Ruthenium  and  Osmium,  410.  Rhodium  and  Iridium,  41 1.  Palladium,  412, 
Platinum,  413. 

Spectrum  Analysis,  415. 


INDEX,  421. 


A TEXT-BOOK 


OF 

INORGANIC  CHEMISTRY. 


INTRODUCTION. 

The  natural  sciences  are  occupied  with  the  investigation  of  the  innu- 
merable substances  and  changes  by  which  we  are  surrounded.  Physics 
and  chemistry  differ  in  a sense  from  the  other  sciences,  in  that  their  domain 
of  research  is  not  restricted  to  any  very  definite  province  of  nature; 
indeed,  it  is  not  confined  to  any  one  planet.  This  is  due  to  the  fact  that 
all  changes,  so  far  as  they  are  perceptible  to  our  senses,  are  referable  to 
chemical  and  physical  causes.  The  best  and  mist  satisfactory  informa- 
tion relative  to  the  ]3roperties  and  composition  of  a substance,  no  matter 
what  its  source,  is  afforded  by  jihysics  and  chemistry.  For  these  reasons, 
therefore,  these  two  sciences — chemistry  and  physics — are  regarded  as  the 
general  natural  sciences  in  contradistinction  to  the  other  more  special 
sciences. 

Physics  deals  with  the  doctrine  of  equilibrium  and  with  that  of 
motions.  The  latter  are  visible,  as  those  of  mass — in  fall,  projection, 
rotation,  propagation  in  a plane,  etc.  ; or  they  are  invisible,  and  are  only 
perceptible  by  their  results — sound,  heat,  light,  electricity.  Chemistry, 
on  the  other  hand,  reveals  to  us  the  composition  of  matter,  and,  in  the 
formation  of  new  compounds,  acquaints  us  with  the  rules  and  laws  by 
which  its  various  forms  act  upon  each  other.  The  domain  of  chemistry 
and  physics  consequently  extends  throughout  all  the  natural  kingdoms, 
and  each  of  the  special  natural  sciences,  even  astronomy,  avails  itself 
of  the  aid  given  by  physics  and  chemistry  to  attain  its  own  particular 
goal.  However,  the.se  two  sciences  are  mutually  dependent  upon  one 
another  for,  so  far  as  we  know,  there  cannot  be  motion  without  matter, 
nor  matter  without  motion. 

The  influence  exerted  by  chemi.stry  and  physic.s  upon  civilization  corresponds  to 
their  exalted  position  in  the  group  of  natural  sciences.  By  means  of  these  sister 
sciences  the  y^roducts  and  forces  of  nature  have  been  more  completely  utilized  than 
ever  before,  hence  there  are  but  few  fields  of  human  activity  which  in  the  course  of  the 
pre.sent  century  have  not  been  enriched  by  the  accumulation  of  chemical  and  physical 
observations.  If  the  conquests  of  chemical  and  physical  investigation  of  even  the  last 
2 17 


cS 


INORCIANIC  CHEMISTRY. 


five  years  were  suddenly  to  disappear  an  almost  unbearable  retrogression  would  be  fell  in 
commerce,  manufacture,  the  various  industries  and  in  agriculture,  and  every  individual 
would  find  himself  having  recourse  to  innumerable  advantages  and  facilities  of  wliic  li  at 
present  he  is  scarcely  conscious. 

The  following  figures  give  some  idea  of  the  importance  of  ( ierman  chemical  industries 
In  1895,  114,581  operators  were  engaged  in  5974  factories  devoted  to  chemical  manu 
facture  ; these  represented  in  round  numbers  100,000,000  M.  or  ^25,000,000  in  wages, 
etc.  The  exports  of  chemical  i)roducts  for  the  same  year  were  valued  at  290,000,000  M. 
or  $72,500,000. 

Compare  : Wickelhaus,  Wirthschaftliche  Hedeutung  chemischcr  Arbeit,  lierlin,  1893  ; 
Ferd,  Fischer,  Das  Studium  der  technischen  Chemie,  Jlraunschweig,  1897. 


A closer  scrutiny  of  natural  olijects  discloses  the  fact  that  they  in 
time  succumb  to  many  more  or  less  serious  alterations  or  changes. 
Although  no  abrupt  boundaries  are  [tresented  in  nature,  but  gradual 
transitions  and  intermediate  steps  throughout,  two  tolerably  distinct 
classes  of  phenomena  may  be  observed.  Some  changes  in  the  condition  of 
bodies  are  only  superficial  (external),  and  are  not  accompanied  by  material 
alteration  in  sid'stance.  Thus  heat  converts  water  into  steam,  which  upon 
subsequent  cooling  is  again  condensed  to  water,  and  at  lower  temperatures 
becomes  ice.  In  these  three  conditions,  the  solid,  li(|uid,  and  gaseous, 
the  substance  or  the  fuatter  of  water  or  ice  is  unchanged;  only  the  sejja- 
ration  and  the  motion  of  the  smallest  particles — their  states  of  aggregation 
— are  different.  If  we  rub  a glass  rod  with  a piece  of  cloth,  the  glass 
acquires  the  property  of  attracting  light  objects,  e.  g.,  particles  of  paper. 
It  becomes  electrified.  An  iron  rod  allowed  to  remain  suspended  verti- 
cally for  some  time  slowly  acquires  the  power  of  attracting  small  pieces  of 
iron.  Through  the  earth’s  magnetism  it  has  become  magnetic.  In  both 
instances  the  glass  and  iron  receive  new  properties;  in  all  other  respects, 
in  their  external  and  internal  form  or  condition,  they  have  suffered  no 
perceptible  alteration;  the  glass  is  glass,  and  the  iron  remains  iron.  All 
such  changes  in  the  condition  of  bodies,  unaccompanied  by  any  real  alter- 
ation in  substance,  are  known  as  physical  phenomena. 

Let  us  now  turn  our  attention  to  the  consideration  of  another  class  of 
phenomena.  It  is  well  known  that  ordinary  iron  undergoes  a change,  which 
we  term  rusting,  i.  e.,  it  is  transformed  into  a brown  substance  which  is  en- 
tirely different  from  iron.  On  mixing  finely  divided  copper  filings  with 
flowers  of  suli)hur  (pulverulent  sulphur)  there  results  an  apparently  uniform, 
grayish-green  powder.  If  this  be  examined,  however,  under  a magni- 
fying glass,  we  can  very  plainly  distinguish  the  red  metallic  copper  par- 
ticles in  it  from  the  yellow  of  the  sulphur;  by  treating  with  water,  the 
sjiecifically  lighter  sulphur  particles  can  easily  be  separated  from  those  of 
tlie  cop])er.  Carbon  bisulphide  will  dissolve  out  the  sulphur  particles. 
Hence  this  powder  represents  nothing  more  than  2l  mechanical  mixture. 
If,  however,  this  mixture  be  heated,  e.  g.,  in  a glass  test-tube,  it  will  com- 
mence to  glow,  and  on  cooling,  a black,  fused  mass  remains,  which  differs 
in  all  res])ects  from  coj)per  and  sulphur,  and  even  under  the  strongest 
microscojie  does  not  reveal  the  slightest  trace  of  the  latter,  and  elutri- 
ation  witli  water  or  treating  with  carbon  bisuli)hide  will  not  affect  a sepa- 
ration of  the  ingredients.  By  the  mutual  action  of  sulphur  and  copper  in 
|)resence  of  heat,  a new  body  with  entirely  different  ])roi)erties  has  been 


INTRODUCTION. 


19 


produced,  and  is  named  copper  sulphide.  Mixtures  of  sulphur  with  iron 
or  with  other  metals  act  in  a similar  manner ; the  resulting  bodies  are 
known  as  sulphides. 

Such  mutual  action  of  different  bodies  occurs  not  only  under  the  influ- 
ence of  heat,  but  frequently  at  ordinary  temperatures.  If,  e.g.,  mercury 
and  sulphur  are  rubbed  together  in  a mortar,  there  is  produced  a uniform, 
black  compound,  called  mercury  sulphide.  The  action  of  gaseous  chlorine 
upon  various  metals  is  quite  energetic.  When  finely  divided  antimony  is 
shaken  into  a flask  filled  with  yellow  chlorine  gas,  flame  is  produced; 
each  antimony  particle  burns  in  the  chlorine  with  a bright  white  light. 
The  product  of  this  action  of  solid  metallic  antimony  and  gaseous  yellow 
chlorine  is  a colorless,  oily  liquid,  known  as  antimony  chloride.  Such 
occurrences,  therefore,  in  which  a complete  and  entire  alteration  takes 
])lace  in  the  bodies  entering  the  reaction,  are  termed  chemical  phenomena. 

In  the  previously  described  experiments  we  observed  the  phenomena 
of  chemical  combination ; from  two  different  bodies  arose  new  homo- 
geneous ones.  The  opposite  may  occur  : the  deco77iposition  of  com- 
pound bodies  into  two  or  more  dissi77iilar  ones.  If  red  mercuric 
oxide  be  heated  in  a test-tube,  it  will  disappear;  a gas  (oxygen)  is 
liberated,  which  will  inflame  a mere  spark  on  wood ; in  addition,  we  find 
deposited  upon  the  upper,  cooler  portions  of  the  tube,  globules  of 
mercury.  From  this  we  observe  that  on  heating  solid  red  mercuric  oxide 
two  different  bodies  arise:  gaseous  oxygen  and  liquid  mercury.  We 
conclude,  then,  that  mercuric  oxide  holds  in  itself,  or  consists  of,  two 
constituents — oxygen  and  mercury.  This  conclusion,  arrived  at  by 
decomposition,  or  analysis.,  may  be  readily  verified  by  combination  or 
synthesis.  It  is  only  necessary  to  heat  mercury  for  some  time,  at  a some- 
what lower  temperature  than  in  the  preceding  experiment,  in  an  atmos- 
])here  of  oxygen,  to  have  it  absorb  the  latter  and  yield  the  compound  we 
first  used — red  mercuric  oxide.  The  direct  decomposition  of  a compound 
body  into  its  constituents  by  mere  heat  does  not  often  happen.  Generally, 
the  cooperation  of  a second  substance  is  required,  which  will  combine 
with  one  of  the  constituents  and  set  the  other  free.  In  this  manner  we 
can,  for  example,  effect  the  decomposition  of  the  previously  synthesized 
mercury  sulphide,  viz.,  by  heating  it  with  iron  filings;  the  iron  unites 
with  the  sulphur  of  the  mercury  sulphide,  to  form  iron  sulphide,  while  the 
mercury  is  set  free. 

If,  in  a similar  manner,  natural  objects  be  decomposed,  bodies  or  sub- 
stances are  finally  reached  which  have  withstood  all  attempts  to  bring 
about  their  division  into  further  constituents,  and  which  cannot  be  formed 
by  the  union  of  others.  Such  substances  are  che7nical  ele77ienis ; they 
cannot  be  converted  into  one  another,  but  constitute,  as  it  were,  the  limit 
of  chemical  change.  Their  number,  at  present,  is  about  70;  some  have 
been  only  recently  discovered.  To  them  belong  all  the  metals,  of  which 
iron,  copper,  lead,  silver,  and  gold  are  examples.  Other  elements  do  not 
possess  a metallic  appearance,  and  are  known  as  7netalloids  (from  I 
resemble).  It  would  be  more  correct  to  term  them  non-7netals.  To  these 
belong  sulphur,  carbon,  phosphorus,  oxygen,  etc.  The  line  between 
metals  and  non-metals  is  not  very  marked.  Thus,  mercury,  despite  the 


20 


INORGANIC  CHEMIS’IRY. 


fad  that  it  is  liquid  at  tlie  ordinary  tcini)craturc,  must  1)C  included  among 
the  metals  because  of  its  chemical  i)roj)ertics. 

All  the  substances  known  to  us  arc  made  up  of  these  elements.  Water 
is  acomj)onnd  of  two  gaseous  elements — hydrogen  and  oxygen  ; common 
salt  consists  of  the  metal  sodium  and  the  gas  chlorine,  'bhe  elements 
make  up  not  only  our  own  earth,  but  the  heavenly  bodies  are  comijosed 
of  them  ; at  least  as  far  as  has  been  proved  by  spectrum  analysis. 


THE  PRINCIPLE  OF  THE  INDESTRUCTIBILITY  OF  MATTER. 

If  the  total  weight  of  substances,  wliich  are  to  act  chemically  upon  one 
another,  be  determined,  and  the  chemical  action  be  then  allowed  to 
occur,  it  will  be  discovered  upon  ascertaining  the  weight  of  the  resulting 
bodies,  if  due  consideration  be  given  for  unavoidable  errors  of  ex})eri- 
ment,  that  no  loss  or  increase  in  weight  has  occurred— no  change  in 
mass,  because  mass  and  weight  are  strictly  i)roj)ortional  to  each  other  for 
one  and  the  same  j)lace.  Jt  is  in  these  cases  immaterial  whether  a com- 
pound body  be  resolved  into  its  elements,  or  whether  elements  unite  to 
produce  com])ound  bodies;  the  })roducts  ])resent  after  the  chemical 
reaction  will  always  weigh  exactly  what  the  bodies  ])receding  the  reaction 
weighed  (comi:)are  the  ex])eriments  of  Landolt,  Ber.  26  (1893),  1820). 
Well-known,  general  phenomena  apparently  contradict  this  scientific 
conclusion.  We  observe  plants  springing  from  a small  germ  and  con- 
stantly acquiring  weight  and  volume.  This  spontaneous  increase  of  sub- 
stance, however,  is  only  seeming.  Closer  inspection  proves  conclusixely 
that  the  growth  of  plants  occurs  only  in  consequence  of  the  alxsorjjtion 
of  substance  from  the  earth  and  atmosphere.  The  opi)osite  phenomenon 
is  seen  in  the  burning  of  combustible  substances,  where  an  apparent 
annihilation  of  matter  takes  place.  But  even  in  this,  careful  observation 
will  discover  that  the  combustion  phenomena  consist  purely  in  a trans 
formation  of  visible  solid  or  liquid  bodies  into  non-visible  gases.  Carbon 
and  hydrogen,  the  usual  constituents  of  combustible  substances,  e.  g., 
petroleum  or  wood,  combine  in  their  combustion  with  the  oxygen  of 
the  air  and  yield  gaseous  products — the  so-called  carbon  dioxide  and 
water — which  diffuse  into  the  atmosphere.  If  these  products  be  col- 
lected, their  weight  will  be  found  not  less,  but  indeed  greater,  than  that 
of  the  consumed  body,  and  this  is  explained  by  the  fact  that  in  addition 
to  the  original  weight  they  have  had  the  oxygen  of  the  air  added.  Such 
a combustion  must,  therefore,  be  regarded  as  a conversion  of  visible 
solid  or  liquid  substances  into  invisible  gaseous  matter. 

d'he  production  (creation)  or  annihilation  of  matter  has  never  been 
demonstrated  as  occurring  in  any  change.  A conq^ound  body  is  com- 
prised of  certain  elements,  and  contains  a very  definite  quantity  by  weight 
of  each  of  them.  If  it  decomixise  it  naturally  breaks  down  into  its  con- 
stituents, which  iierhaps  reunite  in  some  other  manner  to  form  new  com- 
pounds, alwavs,  however,  pres'^rving  their  original  nature,  their  original 
weight  and  their  masses.  This  fundnmcntal  truth  is  ihe  law  of  the  inde- 
stnictihilify  of  matter.  'The  early  (Irecian  ])hilosophers  arrived  at  this 
conclusion  by  a keen  observation  of  the  changes  of  daily  life  (compare 


INTRODUCTION. 


21 


Lucretius — ‘‘Nature”),  and  since  that  time  it  has  always  been  viewed 
as  an  established  fundamental  principle  of  exact,  scientific  investigation. 
Consult  Debus,  Ueber  einige  Fundamental-Satze  der  Chemie,  Kassel,  1894. 

THE  PRINCIPLE  OF  THE  CONSERVATION  OF  ENERGY. 

All  imaginable,  appreciable,  natural  phenomena  have  their  causes,  which 
escape  simple  observation,  and  are  only  realized  by  scientific  research. 
We  observe  that  iron  rusts  on  exposure  to  the  air.  By  chemical  investi- 
gation we  learn  that  rust  is  a compound  of  iron  and  oxygen.  This  latter, 
as  we  also  learn  from  chemistry,  is  a constituent  of  the  air.  If  it  be  re- 
moved from  the  latter,  the  iron  will  cease  to  rust.  Hence  the  tendency 
of  iron  and  oxygen  to  combine  with  one  another  is  designated  as  the  cause 
of  the  rusting.  It  is  generally  said  that  a force  exists  between  them, 
which  effects  their  union  and  this  force  is  called  chemical  affinity.  In 
changes  of  other  kinds  our  ex[)lanation  assumes  the  presence  of  some 
force.  The  falling  of  bodies  is  attributed  to  the  force  of  gravitation — 
a universal  attracting  force,  which  even  influences  the  course  of  the  stars. 
The  decomposition  of  a body  by  external  efforts  acts  in  opposition  to  its 
cohesive  force.  Two  different  bodies  directly  in  contact  with  one  another 
are  held  together  by  the  adhesive  force  existing  between  them.  By  assump- 
tions similar  to  these  it  is  possible  to  refer  an  almost  infinite  number  of 
changes  to  but  comparatively  few  causes.  But,  in  doing  this,  it  must  be 
remembered  that  the  nature  of  the  force  continues  enigmatical,  even  if 
we  know  the  laws,  by  which  it  acts,  as  well  as  we  know  those  of  gravity. 

At  present  the  movements  of  the  ])arts  of  matter  are  considered  to  be 
the  cause  of  other  phenomena  which  formerly  were  ascribed  to  the  actions 
of  special  forces.  It  has  long  been  known  that  a sounding  body  com- 
pletes certain  vibrations,  which  are  imparted  to  the  surrounding  air  and 
arrive  thereby  to  the  tympana  of  our  ears.  The  same  is  true  of  the  phe- 
nomena of  light  and  heat.  It  is  supposed  that  the  heat  phenomena  are 
due  to  an  energetic  oscillating  motion  of  the  smallest  particles  of  a body. 
Upon  grasping  a warm  substance,  motion  present  in  it  is  partially  trans- 
ferred to  us  and  we  experience  warmth.  If,  on  the  contrary,  heat  motion 
be  passed  from  the  hand  to  the  object  touched,  the  latter  appears  to  us  to  be 
cold.  At  sufficiently  elevated  temperatures,  and  indeed  even  under  other 
conditions  (phosi)horescence,  fluorescence)  the  heat  motions  of  bodies 
produce  in  the  surrounding  aether  (a  hypothetical  medium,  enigmatical 
in  its  nature  and  capable  of  penetrating  everything)  motion  extending 
in  all  directions  in  a wave-like  manner  and  known  as  radiant  heat.  If 
these  waves  follow  each  other  rapidly  enough  they  become  light  waves 
and  are  perceived  as  such  by  the  retina  of  the  eye.  The  source  of  electric 
phenomena  and  of  the  x or  Rontgen  rays  is  also  supposed  to  be  due  to 
motion,  of  an  unexplained  nature,  present  in  the  aether.  H.  Hertz  in 
1888  demonstrated  by  the  radiation  of  a periodic  electric  force  reversing 
its  direction  that  these  rays  were  subject  to  the  same  laws  as  those  of  light 
and  radiant  heat.  They  are  not  only  propagated  in  direct  lines  with  all 
the  velocity  of  light  but  are  similarlv  reflected  and  refracted.  Conse- 
quently, in  the  periodic  change  of  direction  we  find  the  same  motion 


22 


INORGANIC  CHEMISTRY. 


alterations  of  the  aether  to  be  at  the  basis  of  liglU  and  radiant  heat,  as 
they  are  the  cause  of  electric  phenomena. 

Numerous  physical  investigations  have  demonstrated  that  the  various 
forms  of  motion  can  be  transferred  mjt  only  from  one  body  to  another 
but  that  they  can  be  converted  into  one  another.  The  bullet  which  in  its 
flight  is  sustained  by  resistance  becomes  hot ; the  visible  motion  of  the 
entire  mass  ceases  and  is  then  converted  into  the  invisible  motion  of  the 
minutest  particles  perceptible  to  us  as  heat.  Conversely,  the  motion  of 
the  smallest  particles  is  changed  to  that  of  large  masses,  if  by  means 
of  the  steam  engine  we  ])roduce  driving  force  by  heat,  and  the  latter  by 
combustion,  a chemical  change. 

It  will  be  discovered  upon  attempting  to  measure  these  changes  that 
the  different  forces  or  modes  of  motion  bear  a fixed  ratio  of  trans- 
formation to  one  another.  According  to  the  proposition,  causa  cequat 
cffectum,  they  are  subject  to  equivalent  convertibility.  At  present  the 
indestructible  portion  of  all  these  changes  is  termed  energy,  which  mani- 
fests itself  as  (i)  jnechanical  energy,  (2)  thermal  energy,  (3)  electric  and 
jna;netic  energy,  (4)  chetnical  and  internal  e7iergy  and  (5)  radiant  energy. 

The  power  of  a body  to  do  work  is  called  energy.  It  is  distinguished  as  potential  and 
kinetic.  Examples  are  as  follows  : The  mass  ?n,  acted  upon  by  the  constant  force  p, 
traverses  the  distance  J and  acquires  the  speed  v.  This  would  be  represented  by  the 

equation  ps  = As  the  force  p has  acted  through  the  distance  s,  it  has  performed 

the  work  ps.  If  a hammer,  weighing  p,  be  raised  to  the  height  s,  the  work  done, 
opposed  to  gravity,  equals /s-.  This  work  is  present  as  potential  energy  in  the  raised 
hamirrer,  so  far  as  the  presence  of  a force  and  its  removal  through  space  can  possibly  make 
a mechanical  action  present.  If  the  hammer  be  allowed  to  fall  it  acquires  kinetic  energy 

-^^>2  (active  force),  which  is  equal  to  the  work  expended  in  raising  the  hammer,  or  the 

potential  energy  of  the  raised  hammer.  In  the  fall  the  potential  is  transformed  into 
kinetic  energy  [Kiveu — to  move  ; kvepyeu — to  act). 

A given  mass-motion  can  be  converted  into  a definite  heat-quantity. 
By  the  application  of  the  latter,  work  can  be  performed  again  equivalent 
to  the  mass-motion  (first  principle  of  the  mechanical  theory  of  heat). 
The  quantity  of  heat  sufficient  to  raise  i kilogram  of  water  at  the  ordinary 
temperature  1°  C.  is  taken  as  the  unit  in  the  determination  of  heat  and 
is  called  calorie  (large).  To  produce  this  unit  by  mass-motion  would 
require  mechanical  work  (at  the  average  geographical  latitude)  equal  to 
426  kilogram-meters.  This  quantity  signifies  that  if  under  the  condi- 
tions mentioned  as  to  plate  i kilogram  falls  through  426  meters,  or  426 
kilograms  through  i meter,  and  the  resulting  total  kinetic  energy  be 
transformed  into  heat,  the  latter  will  raise  i kilogram  of  water  from  15° 
to  16°  in  temperature.  In  the  conversion  of  heat  into  mass-motion, 
however,  one  calorie  will  disappear  for  every  426  kilogram-meters  of 
work  performed.  This  magnitude  (or  constant)  is  known  as  the  tne- 
chanical  equivalent  of  heat. 

Chemical  energy  can  be  measured  by  tlie  heat  or  electricity  developed  in  chemical 
changes.  While  all  forms  of  energy  can  be  converted  without  difficulty  into  heat,  this 
particular  form  is  only  altered  with  limitations  into  the  other  forms.  It  is  only  when 
lieat  passes  from  a warmer  to  a colder  substance  that  a tlelinite  portion  of  it  can  by  proper 


INTRODUCTION. 


23 


appliances  be  changed  into  mechanical  work.  Could  this  so  occur  that  the  greatest  possible 
portion  of  the  heat  could  be  utilized  in  the  performance  of  mechanical  work,  then  the 
process  might  be  reversed.  If,  however,  the  greatest  possible  quantity  of  heat  is  not  con- 
sumed in  doing  such  work  but  is  lost  by  conduction  and  radiation  from  a higher  to  lower 
temperature,  then  the  change  is  not  reversible.  For  a portion,  at  least,  of  the  heat  there 
is  no  possibility  of  change  to  mechanical  work  ; it  is  permanently  lost  {degradation  of 
energy).  The  reversibility  is  made  impossible  by  such  changes  (conduction,  radiation,  fric- 
tion, etc.),  which  take  place  under  given  conditions,  of  their  own  accord.  As  the  op- 
posite conversions  do  not  occur  spontaneously  (the  passage  of  heat  from  a body  of 
lower  to  one  of  higher  temperature),  in  nature  the  changes  of  the  first  order  must  exceed 
those  of  the  second  order.  The  sum  of  all  the  conversion-values  (calculated  as  proceeding 
positively  in  the  sense  of  the  first  group)  is  termed  entropy  {rpoirri — conversion).  The 
entropy  of  all  nature  is  therefore  in  the  act  of  constant  increase  (Clausius,  Second  Principle 
of  the  Mechanical  Theory  of  Heat). 

The  law  of  the  conservation  of  energy,  according  to  which  the  energy 
in  nature  is  of  a convertible  character  but  unalterable  in  quantity,  con- 
stitutes one  of  the  most  important  foundations  of  the  science  of  nature. 

It  was  first  clearly  explained  and  definitely  enunciated  by  Julius  Robert  Mayer  (1842), 
a physician  of  Heilbronn  (Annalen,  42,  233  ; see  also  his  “ Mechanik  der  Warme  gesam- 
melte  Schriften  ”),  Not  knowing  of  Mayer’s  work,  or  of  the  treatise  of  Colding  (1843), 
a Dane,  in  which  the  principle  of  energy  was  also  developed,  Hermann  v.  Helmholtz 
(1847)  announced  the  law  as  empirical,  followed  and  developed  it  mathematically 
through  all  the  domains  of  natural  phenomena  (Ueber  die  Erhaltung  der  Kraft,  Berlin, 
1847).  Mayer  was  also  the  first  to  discover  the  mechanical  equivalent  of  heat,  which 
shortly  afterwards  James  Prescott  Joule  definitely  determined  by  accurate  experiments. 

Heat  almost  invariably  appears  in  chemical  union  ; even  light  and 
electricity  can  be  produced  by  chemical  processes,  or  work  can  be  per- 
formed in  opposition  to  external  pressure  by  increase  in  volume.  All 
these  forms  of  energy  owe  their  origin  to  the  potential  energy  of  the 
chemical  forces,  which  in  the  process  of  chemical  change  do  work. 
Hence,  we  may  speak  of  chemical  energy  or  of  chemical  tension.  In  the 
chemical  decomposition  of  a compound  body  into  its  components,  on  the 
other  hand,  heat  is  usually  absorbed  and  disappears  as  such  and  becomes 
chemical  energy.  Thus,  in  the  union  of  approximately  i kilogram  of 
hydrogen  with  8 kilograms  of  oxygen  to  produce  9 kilograms  of  water 
a quantity  of  heat,  equivalent  to  34,200  calories,  is  set  free,  and  this 
corresponds  to  work  equivalent  to  34,200  X 426  = 14,569,200  kilogram- 
meters.  In  the  decomposition,  on  the  other  hand,  of  9 kilograms  of  water 
into  hydrogen  and  oxygen,  the  same  force  or  quantity  of  heat  is  neces- 
sary. Therefore,  the  same  quantity  of  force  or  motion  must  be  contained 
in  the  form  of  chemical  energy  in  the  liberated  hydrogen  and  oxygen. 

Although  all  bodies,  and  the  elements  especially,  possess  chemical 
energy,  they  do  not  manifest  it  in  the  same  way.  Some  of  them  react 
readily  with  one  another  and  others  with  difficulty  or  not  at  all.  The 
cause  of  this  variation  in  behavior  is  entirely  unknown  to  us.  It  is  cus- 
tomary to  express  the  fact  by  saying  the  bodies  have  a strong,  a feeble,  or 
no  affinity  for  one  another.  Formerly,  bodies  which  combined  chem- 
ically, were  supposed  to  be  related  to  each  other,  and  it  was  assumed  that 
their  affinity — their  tendency  toward  one  another — was  satisfied  by  their 
union.  This  choice  of  terms  was  unfortunate.  The  nature  of  the  chemical 
attraction,  which  produces  and  holds  chemical  compounds  together  is  just 


24 


INORGANIC  CHEMISTRY. 


as  enigmatical  as  that  of  gravity.  Even  the  laws  in  accordance  with 
which  affinity  acts  are  scarcely  known  to  us.  It  is  only  in  recent  times 
that  advances  have  been  made  in  this  direction,  since  the  positions  and 
the  motions  of  atoms  and  of  molecules  have  been  regarded  as  affording 
hints  as  to  the  cause  of  a reaction. 

CONSTITUTION  OF  MATTER.  ^ 

>- 

ATOM  AND  MOLECULE. 

If  we  seek  to  give  expression  to  the  constitution  of  the  chemical  ele- 
ments and  the  bodies  composed  of  them  we  return,  if  guided  by  experience, 
to  the  ancient  atomic  hypothesis  which  alone  is  justified  by  the  present 
condition  of  chemical  and  ])hysical  investigation.  It  apjiears  that  the 
Indian  and  Grecian  natural  philosophers  established  this  same  hypothesis 
in  a purely  inductive  manner  [Kanada  (founder  of  the  Vaiseshika  system), 
Lucippus,  Democritus  (500  b.  c.)  and  Ei)icurus  (400  b.  c.)]. 

The  foundation  of  this  hypothesis  is  brilliantly  set  forth  in  the  first  book  of  Lucre- 
tius (died  55  B.  c. ) “ Ueber  die  Natur  der  Dinge.”  Numerous  observations  are  re- 
corded. The  following  occurs  there  in  substance  : All  bodies  can  be  divided  into  in- 
finitely small  parts,  no  longer  recognizable  by  sight  or  taste.  Invisible  aqueous  vapor 
separates  from  the  air  as  water  upon  cold  objects  and  again  disappears  on  the  ap])roach  of 
heat.  The  ring  that  is  constantly  worn  on  the  finger  becomes  thinner  in  the  course  of 
years.  Dropping  water  wears  away  the  stone.  The  smooth  pavement  is  made  rough  by 
walking.  All  these  things  occur  without  our  perceiving  what  at  any  one  time  departs 
from  the  ring,  etc.  Hence  we  conclude  the  bodies  are  composed  of  invisible,  extremely 
small  parts,  which  to  us  are  without  mass.  These  particles,  the  atoms  (arofioQ — indivisible), 
are  indestructible  and  cannot  be  created.  There  is  nothing  beyond  them  and  the  vacant 
space  between.  The  difference  in  things  is  due  to  the  difference  in  number,  size,  form, 
and  arrangement  of  the  atoms.  There  is  no  qualitative  difference  in  the  atoms,  they  act 
upon  one  another  only  by  contact  and  pressure.  Change  is  only  a union  and  separation 
of  atoms  ; nothing  occurs  by  chance  but  everything  by  reason  and  with  necessity.  Thus 
far  Lucretius. 

In  the  first  half  of  the  17th  century  the  atomic  idea,  which  until  then  had  been 
driven  into  the  background  by  the  Aristotelian  philosophy,  was  resuscitated  by  Daniel 
Sennert,  a German  physician,  and  a P'rench  ecclesiastic  named  Pierre  Gassendi.  They 
adopted  the  Greek  atomic  doctrine  ; and,  from  the  point  of  view  of  the  atomists,  constitute 
the  connecting  link  between  the  past  and  the  present.  Since,  however,  our  modern 
atomic  notions  have  developed  step  by  step  from  the  ideas  of  Sennert  and  Gassendi,  their 
beginnings  go  back  to  and  have  their  origin  in  Lucippus  and  Democritus.  Atoms  were 
introduced  into  chemistry  by  Robert  Boyle,  a contemporary  and  follower  of  Gassendi. 
Boyle  was  the  first  chemist  to  devote  his  experiments  to  the  noble  purpose  of  investi- 
gating nature.  (Compare  F.  A.  Lange,  Geschichte  des  Materialismus,  3 Aufl.  1876  ; 
Grie.sbach  physik.-chem.  Propadeutik  (1895),  and  the  fasciculus  of  Debus  referred  to 
on  page  2I.) 

'The  scientific  foundations  of  our  present  atomic  doctrine  will  be 
described  later.  Its  fundamental  ideas  alone  will  be  given  now.  We 
assume  that  an  element  consists  of  atoms  perfectly  similar  to  one  another, 
but  differing  from  those  of  other  elements.  We  must  grant  that  there  are 
as  many  kinds  of  atoms  as  there  are  different  elements.  A compound  body 
like  iron  sulphide,  according  to  this  view,  is  produced  by  the  combination 
of  sulj)hur  atoms  with  iron  atoms  in  a definite  ratio.  Those  ])articles  of 
a compound,  representing  the  limit  of  divisibility  so  far  as  similarity 


. INTRODUCTION. 


25 


goes,  are  called  molecules  (j?iolecula,  diminutive  of  moles — the  mass),  by 
further  division  they  are  resolved  into  dissimilar  parts.  Hence,  iron 
sulphide  is  made  up  of  molecules,  which  in  turn  consist  of  atoms  of  iron 
and  of  sulphur.  We  shall  learn  later  that  the  elements — with  few  excep- 
tions—are  composed,  at  ordinary  temperature,  not  of  a collection  of  free 
atoms,  but  of  an  aggregation  of  atom-groups — of  molecules.  A molecular 
structure  is  the  rule.  In  the  case  of  the  elements  the  molecules  are  com- 
posed of  like  kinds  of  atoms,  in  compounds  they  consist  of  dissimilar 
atoms. 

The  greatest  advancement  of  the  atomic  theory  is  due  to  John  Dalton. 
By  assuming  that  the  atoms  combined  with  one  another  in  definite  pro- 
portions, he  laid  the  basis  for  the  determination  of  the  relative  atomic 
weights,  and  thereby  became  the  founder  of  the  chemical  atomic  theory 
based  upon  definite  weight  proportions  (1804).  If  two  elements  form 
but  a single  compound  by  union  with  one  another,  it  may  be  assumed 
with  Dalton,  as  long  as  no  other  reason  to  the  contrary  exists,  that  their 
molecules  consist  of  an  atom  of  each  of  the  two  elements.  Should  two 
compounds  of  the  elements  A and  B be  known,  then  the  molecule  of  the 
one  compound  would  consist  of  an  atom  each  of  A and  of  B,  consequently 
of  two  atoms,  while  the  other  compound  might  be  composed  of  three 
atoms  (2  A -f  B or  2 B -f  A),  etc.  With  these  premises  clearly  enunci- 
ated it  is  possible  to  determine  the  relative  atomic  weights  of  the  elements. 

One  hundred  parts  of  the  previously  mentioned  iron  sulphide  consist 
in  round  numbers  of  63.6  parts  of  iron  and  36.4  parts  of  sulphur.  If, 
however,  in  accordance  with  the  assumptions  of  Dalton,  there  is  in  this 
compound  one  atom  of  sulphur  for  one  atom  of  iron,  then  the  atomic 
weights  of  iron  and  sulphur  must  be  to  one  another  as  63.6  : 36.4.  The 
ratios  between  the  atomic  weights  of  the  elements  may  be  determined  in 
this  manner.  If  for  any  element  a number  be  taken  for  its  atomic 
weight,  it  can  readily  be  calculated  in  what  ratio  the  atomic  weights  of 
all  the  other  elements  stand  to  this  arbitrarily  chosen  standard,  and  we 
thus  obtain  the  relative  ato77iic  weights.  That  element  which  combines 
with  the  majority  of  the  other  elements  to  form  compounds  capable  of 
the  most  accurate  analysis,  is  chosen  as  the  standard  of  comparison. 
Finally,  the  number  or  value  assigned  this  standard  element  as  its  atomic 
weight  is  a matter  of  consensus  of  opinion. 

Later  we  shall  become  acquainted  with  physico-chemical  methods  of 
testing  the  atomic  numbers,  derived  in  a chemical  way,  and  especially  for 
establishing  whether  it  is  not  a fraction  or  a multiple  of  the  true,  relative 
atomic  weight. 

CHEMICAL  SYMBOLS  AND  FORMULAS. 

The  chemical  elements  are  simjfiy  and  conveniently  represented  by  the 
initials  of  their  Latin  or  Greek  names  in  accordance  with  the  suggestion 
of  the  great  Swedish  chemist,  John  Jacob  Berzelius  (1779-1848)  to 
whom  we  are  also  indebted  for  the  first  accurate  atomic  weight  determina- 
tions. Thus  hydrogen  is  designated  by  the  letter  H,  from  the  word 
hydrogenium  ; nitrogen  by  N,  from  nitrogenium.  Elements  having  the 
same  initials  are  distinguished  by  adding  a second  letter;  thus,  Na  indi- 
3 


26 


INORGANIC  CHEMISTRY. 


cates  natrium  (sodium),  Ni — nickel,  Hg — mercury  (from  hydrargyrum), 
Pd — palladium,  Pt — platinum,  etc. 

The  following  table  contains  the  names  of  71  known  chemical  ele- 
ments, together  with  their  symbols  and  their  atomic  weights,  refer7'cd  to 
oxygen  equal  to  16. 

It  may  be  added  : 

That  as  a rule  the  atomic  weights  are  given  with  only  as  many  decimals  as  are  accu- 
rate to  the  last  figure.  In  the  case  of  bismuth,  nickel  and  tin,  indicated  by  *,  this  rule 
is  not  adhered  to.  Also  in  the  ca.se  of  hydrogen  the  more  accurate  value  1. 008  is  given 
as  1. 01  for  ordinary  use. 

The  elements  whose  names  have  ? attached  are  in  doubt  either  as  to  their  simplicity, 
or  as  to  entire  units  in  their  atomic  weights. 


Elements. 

j 

Symbol. 

Atomic 

Wekjiit. 

Elements. 

Symbol. 

Atomic 
Weigh  r. 

Aluminium, 

A1 

27. 1 

Neodymium  (?),  . . . 

Nd 

144 

Antimony  (Stibium), 

Sb 

120 

Nickel, 

Ni 

58.7* 

Argon  (?), 

A 

40 

Niobium, 

Nb 

94 

Ai’senic, 

As 

75 

Nitrogen,  

N 

14.04 

Barium, 

Ba 

137-4 

Osmium, 

Os 

191 

Beryllium, 

Be 

9-1 

Oxygen, 

0 

16.00 

Bismuth,  

Bi 

208.5* 

Palladium, 

Phosphorus, 

Pd 

106 

Boron,  

B 

II 

P 

31.0 

Bromine, 

Br 

79.96 

Platinum, 

Pt 

194.8 

Cadmium,  

Cd 

1 12 

Potassium  (Kalium), 

K 

39-15 

Caesium, 

Cs 

133 

Praseodymium  (?),  . . 

Pr 

140 

Calcium,  

Ca 

40 

Rhodium, 

Rubidium, 

Rh 

103.0 

Carbon, 

C 

12.00 

Rb 

854 

Cerium, 

Ce 

.140 

Ruthenium, 

Ru 

101.7 

Chlorine,' 

Cl 

35-45 

Samarium  (?),  . . . . 

Sm 

150 

Chromium,  

Cr 

52.1 

Scandium, 

44-1 

Cobalt, 

Co 

59 

Selenium, 

79-1 

Copper, 

Cu 

63.6 

Silicon, 

Si 

28.4 

Erbium  (?), 

Er 

166 

Silver  (Argentum),  . . 

Ag 

107.93 

Fluorine, 

FI 

19 

Sodium  (Natrium),  . . 

Na 

23-05 

Gallium, 

Ga 

70 

Strontium, 

Sr 

87.6 

Germanium, 

Ge 

72 

Sulphur, 

S 

32.06 

Gold  (Aurum),  .... 

All 

197.2 

Tantalum, 

Ta 

183 

1 leliuin  (?), 

He 

4 

Tellurium,  ...... 

Te 

127 

Hydrogen, 

H 

1. 01 

Thallium, 

T1 

204. 1 

Indium, 

In 

1 14 

Thorium, 

Th 

232 

Iodine, 

I 

126.85 

'bin  (Stannum),  .... 

Sn  1 

118.5* 

Iridium, 

Ir 

193 

Titanium, 

Tungsten  (Wolfram),  • 

I'i 

48. 1 

Iron  (Ferrum),  .... 

Fe 

56.0 

W 

184 

Lanthanum, 

La 

138 

Uranium, 

Ur 

239-5 

Lead  (Plumbum),  . . . 

PI) 

206.9 

Vanadium, 

Yd 

51.2 

I.ithium,  

Li 

7-03 

Ytterbium, 

Yb 

173 

Magnesium, 

Mg 

24.36 

Yttrium, 

Y 

89 

Manganese, 

Mn 

55-0 

Zinc, 

Zn 

65.4 

Mercury, 

Molybdenum,  .... 

Hg 

Mo 

200. 3 
96.0 

Zirconium, 

Zr 

90.6 

In  nddilioi)  to  the  elements  mentioned  in  the  table  the  following  are  believed  to  have 
been  observed  in  certain  rare  minerals:  terbium  {^niosandrium^  philippiitnt^  7m, 


INTRODUCTION. 


27 


gadolinium^  decipium^  holmhcm,  thuliiun  and  dysprosium.  They  occur  associated  with 
cerium,  lanthanum,  scandium,  ytterbium  and  yttrium.  They  are  so  much  alike  chemic- 
ally that  it  is  difficult  to  separate  them  and  obtain  them  pure  ; they  are  probably  mixtures 
of  unknown  elements. 

Consult  the  paragraphs  under  Air  for  an  account  of  what  seem  to  be  elementary  gases 
(metargon,  neon,  krypton,  xenon)  which  have  been  isolated  from  the  atmosphere.  For 
the  reasons  leading  to  the  adoption  of  oxygen  as  the  element  of  comparison  and  estab- 
lishing its  atomic  weight  as  16,  see  p.  73. 

Compounds  produced  by  the  union  of  the  elements  are  represented  by 
placing  their  corresponding  symbols  together  and  designating  these 
chemical  formulas.  Common  salt,  a compound  of  sodium  and  chlorine, 
is  represented  by  the  formula  NaCl ; mercuric  oxide,  a compound  of 
mercury  and  oxygen,  by  HgO;  iron  sulphide  by  FeS ; hypochlorous 
acid,  a compound  of  hydrogen,  chlorine  and  oxygen,  by  HCIO. 

By  these  premises  chemical  formulas  acquire  a very  precise  and  evident 
importance.  The  formula  NaCl  represents  the  union  of  i atom  of  sodium 
with  I atom  of  chlorine,  and  indicates  that  in  it  23.05  parts,  by  weight, 
of  sodium  are  combined  with  35.45  parts  by  weight  of  chlorine  to  yield 
58.50  parts  of  sodium  chloride  (common  salt).  If  several  atoms  of  an 
element  are  present  in  a compound,  this  is  denoted  by  numbers  which  are 
attached  to  the  symbol  of  the  atom  : 

HCl  H2O  NHg  CH^ 

Hydrochloric  acid.  Water.  Ammonia.  Methane. 

The  formula  of  water  (HgO)  means  that  its  molecule  consists  of  2 
atoms  of  hydrogen  (2.02  parts  by  w^eight)  and  i atom  of  oxygen  (16 
parts  by  weight).  The  formula  of  sulphuric  acid  (H2SOJ  indicates  it  to 
be  a compound  consisting  of  i atom  of  sulphur  (32.06  parts),  4 atoms  of 
oxygen  (4  X 16=  64  parts),  and  2 atoms  of  hydrogen  (2  X i-oi  = 2.02 
parts),  from  which  the  composition  of  the  acid  may  be  at  once  calculated 
into  per  cent.,  or  into  any  desired  quantity  by  weight. 


Atomic  Composition.  In  Per  Cent. 

Hydrogen,  Hg  = 2.02 2.06 

Sulphur,  . . S = 32.06 32.69 

Oxygen,  ..  O4  = 64.00 65.25 

H2SO4  = 98.08 100.00 


A chemical  change  is  represented  by  arranging  these  symbols  in  the 
form  of  an  equation.  The  left  side  of  the  equation  indicates  the  sub- 
stances present  before  the  reaction  occurs,  while  the  right  side  shows  the 
products.  Thus  the  chemical  equation  : 

HgS  -f  Fe  = FeS  -f  Hg 

means  that  mercury  sulphide  (232.3-6 \parts)  and  iron  (56  parts)  have 
combined  to  form  iron  sulphide  (88.06  parts)  and  mercury  (200.3  parts). 
The  equation 

zH  + O = HgO 


indicates  that  i molecule  of  water  has  been  formed  by  the  union  of  2 
atoms  of  hydrogen  with  i atom  of  oxygen. 


28 


INORGANIC  CHEMISTRY. 


It  is  true  these  equations  give  no  idea  of  the  conditions  surrounding 
these  transpositions,  nor  of  the  alterations  in  energy  which  accompany 
tliLMii.  'I'hey  represent  tlie  i)urely  material  side  of  the  change.  'I'hey  indi 
cate  tiie  quantities  by  weight  of  the  substances  entering  the  reaction  aiul 
also  of  the  products:  the  weight  of  the  bodies  entering  the  reaction  is 
equal  to  that  of  the  resulting  })roducts.  Iwery  chemical  equation  is  there- 
fore an  ex])ression  of  the  law  of  the  indestructibility  of  matter  (p.  20). 


CONDITIONS  OF  CHEMICAL  CHANGE. 

THERMO-CHEMICAL  PHENOMENA. 

The  first  requisite  for  bodies  to  act  chemically  iq)on  one  another  is 
that  they  be  brought  into  most  intimate  contact  because  chemical  action 
does  not  occur  at  great  distances.  In  the  case  of  solids  this  intimate 
contact,  so  essential  for  comi)lete  chemical  transformation,  cannot  ordi- 
narily be  attained  by  mere  mechanical  mixing  : the  necessary  condition 
is  best  reached  by  liquefying  the  bodies,  or  at  least  one  of  them,  by 
fusion  or  solution  in  some  solvent.  Hence  the  old  saying  corpora  non 
agiifit  nisi  fluida. 

In  many  instances,  however,  the  chemical  transposition  does  not  take 
place  even  with  the  most  intimate  contact.  An  external  physical  im- 
pulse occasioned  by  light,  by  electricity,  by  change  in  pressure  (Spring, 
van’t  Hoff),  but  more  especially  by  the  temperature,  is  required  for  its 
occurrence.  Thus,  for  example,  hydrogen  and  oxygen  at  the  ordinary 
temperature  are  wholly  indifferent  to  one  another  despite  the  fact  that  as 
gases  they  may  be  mixed  as  completely  as  it  is  possible.  It  is  only  when 
they  are  heated  that  they  combine — slowly  at  200°,  but  with  violent  ex- 
plosion about  700° — to  water.  The  same  occurs  upon  passing  the  electric 
spark  through  the  mixture.  A mixture  of  hydrogen  and  chlorine  will 
remain  unchanged  in  the  dark,  while  in  diffused  sunlight  the  gases  will 
slowly  unite  to  hydrochloric  acid,  but  in  direct  sunlight,  upon  the  appli- 
cation of  heat,  or  by  passing  the  electric  spark  they  will  unite  at  once  with 
great  violence.  A mixture  of  iron  and  mercury  sulphide  requires  the  aid 
of  heat  to  bring  about  the  transposition  to  iron  sulphide  and  mercury. 
At  the  ordinary  temperature  they  appear  to  exercise  no  visible  chemical 
action  upon  one  another.  By  the  application  of  external  energy — heat, 
light,  electricity,  etc. — the  atomic  structure  of  the  molecules  of  hydro- 
gen and  oxygen,  of  chlorine  and  hydrogen,  of  mercury  sulphide  and 
iron,  etc. — is  first  loosened  or  disintegrated  and  then  the  chemical  action 
between  the  several  com})onents  takes  place.  However,  the  experiments 
of  Raoul  Pictet  [1892;  compare  Ber.  26  ( 1893),  iv,  i and  A.  Welter,  Die 
tiefen  'rem])eraturen,  Crefeld  1895]  show  that  chemic  al  reactions  do  not 
occur  at  temjicratures  below — 125°.  Substances  which  at  the  ordinary 
temperature  react  with  the  greatest  readiness — 0.  suljihuric  acid  and 
barium  cliloride,  hydrochloric  acid  and  silver  nitrate,  sodium  and  alcohol 
— appear  at — 80°  to  be  as  indifferent  toward  one  another  as  mercury  sul- 
phide  and  iron,  and  hydrogen  and  oxygen  at  the  ordinary  temperature. 
Even  such  delicate  tests  as  those  of  blue  litmus  and  sulphuric  acid  or  hydro- 


INTRODUCTION. 


29 


chloric  acid  do  not  take  place  below  — 1 10°.  It  is  obvious,  therefore,  that 
the  power  of  substances  to  act  chemically  upon  one  another  is  entirely  de- 
pendent upon  the  external  conditions,  particularly  the  temperature,  pre- 
vailing at  the  moment  of  their  contact.  This  has  been  further  demonstrated 
by  the  experiments  which  it  has  been  possible  to  conduct  at  very  high  tem- 
peratures. It  is  well  known  that  hydrogen  and  oxygen  unite  to  form 
water  above  200°,  but  at  2000°  and  above,  water  breaks  down  again  into 
hydrogen  and  oxygen  (compare  Dissociation  of  Water).  The  higher 
the  temperature  the  more  complete  will  be  the  decomposition,  and 
eventually  a point  will  be  reached  at  which  the  hydrogen  and  oxygen  will 
exert  as  little  chemical  action  upon  one  another  as  they  would  below  200°. 

. A chemical  compound  is,  therefore,  only  wholly  stable  within  a certain 
range  of  temperature.  The  latter  may  change,  however,  with  the  pres- 
sure and  in  the  case  of  some  substances  we  may  not  be  able  to  obtain  it 
with  our  present  facilities.  The  compound  will  begin  to  separate  into 
its  constituents — its  elements — just  as  soon  as  this  limit  is  exceeded.  The 
rapidity  of  the  decomposition,  the  magnitude  to  which  it  may  extend, 
is  also  dependent  upon  pressure  and  temperature.  This  is  therefore  due 
to  the  fact  that  a portion  of  the  product  of  decomposition  reunites  to 
form  the  original  body  and  in  this  way  a state  of  chemical  equilibrium  is 
produced.  The  extent  of  the  decomposition  is  always  definite  for  given 
external  relations.  Similarly,  the  action  of  unlike  bodies  upon  one  an- 
other is  frequently  complete  only  within  definite  ranges  of  temperature, 
and  the  rapidity  with  which  this  action  proceeds  is  in  like  manner  influ- 
enced not  only  by  temperature  but  also  by  pressure  and  by  quantity- 
relations.  Very  often  opposite  reactions,  union  and  decomposition,  oc- 
cur simultaneously,  and  occasion  a state  of  equilibrium. 

Every  chemical  change  is  invariably  accompanied  by  a change  of 
energy — by  the  disengagement  or  absorption  of  heat  (electricity,  etc.). 
The  customary  chemical  equations,  such  as  are  employed  upon  page  27, 
represent  merely  the  material  side  of  a chemical  reaction,  the  nature  and 
quantities  by  weight  of  the  reacting  and  resulting  substances.  But  when, 
for  example,  2.02  grams  of  hydrogen  and  16.00  grams  of  oxygen  unite  to 
yield  water,  there  is  an  accompanying  dynamical  change,  a definite  and 
considerable  quantity  of  heat  is  disengaged — in  this  instance  equaling 
68.4  calories  (p.  23).  An  equation  showing  the  union  of  hydrogen  with 
oxygen  to  form  water — an  equation  which  would  include  both  weight  xtdiC- 
tions  and  those  of  energy — would  read  as  follows: 

2H  + O = HP -f  68.4  Cal. 

Similarly,  in  the  union  of  i.oi  grams  of  hydrogen  with  35.45  grams  of 
chlorine,  forming  36.46  grams  of  hydrogen  chloride,  we  have  a liberation 
of  22  calories : 

H-f-Cl  =HC1  T 22  Cal.  ; 

while  the  formation  of  hydrogen  iodide  (127.86  grams)  from  hydrogen 
(i.oi  grams)  and  iodine  (126.85  grams)  is  accompanied  by  an  absorption 
of  energy — an  absorption  of  6 calories  : 

H -y  I + 6 Cal.  = III. 


30 


INORGANIC  CHEMISTRY. 


But  this  e(|uati()n  is  in  no  manner  to  he  understood  as  meaning  that  l)ydro^en  and 
iodine  reciuire  the  addition  of  l)ut  6 ealories  in  order  that  hydrogen  iodide  may  result. 
The  application  of  that  amount  of  heat  would  only  produee  a warmed  mixture  of  hydro- 
gen and  iodine  vapor.  'The  (juantity  of  energy  e(juivalent  to  6 calories  of  heat  can  only 
be  taken  up  by  the  mixture  under  certain  delinite  physical  conditions  in  such  a way  that 
hydrogen  iodide  is  the  product. 

When  we  desire  to  decompose  water  into  hydrogen  or  oxygen  we  mtist 
restore  all  the  energy  which  has  esca^ted.  This  energy  is  contained  in 
the  free  elements  as  chemical  (potential)  energy.  The  decomposition  of 
hydrogen  iodide  into  its  elements,  on  the  contrary,  occurs  with  a disen- 
gagement of  heat.  This  heat  is  equivalent  to  the  energy  which  the  iodine 
and  hydrogen  took  up  in  their  i)assage  into  hydrogen  iodide. 

Reactions  are  distinguished  as  exothermic  and  endothermic^  i.  e.,  they 
have  positive  or  negative  the7'7?ial  values  depending  upon  whether  heat 
(energy)  is  liberated  or  absorbed. 

The  heat  modulus,  attending  a chemical  reaction,  offers  important  con- 
clusions as  to  the  concurrent  alterations  in  energy,  as  well  as  to  the  nature 
of  the  substances  entering  the  reaction. 

These  relations  are  most  evident  in  those  reactions  in  which  only  two 
ele77ients  participate.  Compounds,  formed  from  their  elements  with  the 
liberation  of  heat,  contain  less  energy  than  the  elements  themselves,  are 
more  stable  than  their  mixture,  and  can  be  resolved  into  the  original 
elements  by  the  consumption  of  energy.  The  conditions  requisite  for 
their  formation  are  equally  present  in  the  mixture  of  the  parent  sub- 
stances, because  every  system  of  bodies,  as  taught  by  mechanics,  strives  to 
attain  that  state  of  equilibrium,  in  which  the  content  of  energy,  convert- 
ible into  work,  of  tension  or,  as  Helmholtz  terms  it,  of  free  energy^  is  as 
low  as  possible.  It  is  for  this  reason,  therefore,  that  reactions  of  this 
description  occur  almost  immediately  on  bringing  the  respective  bodies 
together.  An  example  of  this  kind  has  already  been  given : antimony 
and  chlorine  unite  instantaneously  to  antimony  chloride  (p.  19).  As  a 
rule,  however,  an  external  impulse  from  heat,  electricity,  or  light  is 
needed  to  start  the  reaction.  Several  examples  illustrating  this  have  been 
presented.  This  behavior  must  be  regarded  as  a sort  of  liberation  of  the 
chemical  tension.  It  is  necessary  that  a portion  of  the  molecules — the 
physical  individuals — first  be  resolved  into  their  atoms — the  chemical ‘in- 
dividuals— the  combination  of  the  atoms  in  the  molecules  must  be  made 
less  intimate  at  least.  When  once  the  reaction  has  been  in  this  manner 
begun  at  one  point  of  the  mixture,  it  generally  continues  and  proceeds  of 
itself,  according  to  the  amount  of  heat  developed,  with  greater  or  less 
intensity,  which  may  even  reach  to  explosion,  as  in  the  production  of 
water  or  hydrochloric  acid  from  their  elements. 

Com])ounds  which  absorb  heat  when  produced  from  their  elements,  e.  g., 
hydrogen  iodide  and  nitrogen  chloride,  contain  more  free  energy,  more 
tension,  than  the  parent  sul)stances,  and  are  consequently  less  stable  than 
their  mixture,  d’hey  cannot  be  formed  from  their  elements  without  the 
simultaneous  addil ion  of  energy.  In  this  instance  to  simply  initiate  the 
reaction  is  not  sufficient,  for  energy  must  be  continually  added,  otherwise 
the  chemical  action  will  cease. 


INTRODUCTION. 


31 


Viewed  thermally,  the  decomposition  of  compounds  into  their  elements 
proceeds  oppositely  to  their  formation.  If  the  latter  was  accompanied 
by  heat  evolution  then  the  decomposition  would  occur  with  heat  (energy) 
absorption,  would  require  the  constant  addition  of  energy,  advance  very 
gradually,  never  in  an  explosive  manner,  and  would  be  circumscribed  by 
the  opposing  combination-tendency  of  the  products  of  the  decomposition 
(compare  Dissociation  of  Water).  Just  as  every  transition  of  a system 
from  a condition  of  stability  to  one  of  less  stability  requires  an  expen- 
diture of  work,  so  does  the  decomposition  of  such  a body  as  indicated  in 
the  preceding  lines.  Compounds,  however,  which  like  hydrogen  iodide 
and  nitrogen  chloride  are  produced  with  heat  absorption,  generally  break 
down  with  ease  and  completely  into  their  elements.  In  this  way  the 
system  passes  into  a more  stable  condition.  Frequently,  an  external  im- 
pulse is  all  that  is  necessary  to  start  such  a decomposition.  It  then  pro- 
ceeds of  its  own  accord  and  may  increase  even  to  explosion.  Many  of 
the  bodies  belonging  in  this  group  are  explosive  ; indeed,  some  of  them, 
e.  g.,  chlorine  monoxide  and  nitrogen  iodide  explode  when  touched  or 
when  warmed,  others  again  require  a more  energetic  concussion.  Thus, 
nitric  oxide,  acetylene,  and  cyanogen  explode  if  a slight  amount  of 
mercury  fulminate  be  ignited  in  them. 

The  heat  modulus  is  not  a measure  of  the  affinity  of  the  elements  which 
combine  with  one  another.  Even  the  formation  of  water  from  hydrogen 
and  oxygen,  an  apparently  very  simple  process,  is  really  the  product  of  a 
number  of  chemical  and  physical  changes,  proceeding  by  grades,  which  in 
turn  are  accompanied  partly  by  positive  and  partly  by  negative  thermal 
values— breaking  down  of  the  molecules  into  atoms,  union  of  the  differ- 
ent atoms  to  molecules;  diminution  of  the  number  of  molecules,  lique- 
faction of  the  aqueous  vapor.  The  quantity  of  heat  observed  merely 
represents  the  algebraic  sum  of  all  these  heat  moduli  (thermal  values). 

If  more  than  two  elements,  if  several  compound  substances,  take  part 
in  a chemical  change  the  meaning  of  the  thermal  value  accompanying  it 
is  more  difficult  to  comprehend.  The  majority  of  reactions  of  this  kind 
proceed  in  harmony  with  the  principle  of  greatest  heat  development  3.cc.or&- 
ing  to  which  from  a given  system  of  bodies,  without  the  introduction  of 
external  energy,  that  new  compound  will  result,  in  whose  formation  there 
is  the  greatest  heat  development  (Berthelot).  This  principle  is  not  uni- 
versally acknowledged,  neither  do  all  the  facts  support  it,  nor  is  it  justified 
from  the  standpoint  of  the  mechanical  theory  of  heat.  The  entropy- 
principle  is  a more  acceptable  substitute  (p.  23).  When  the  thermal 
relations  of  the  various  groups  of  elements  and  compounds  are  discussed 
this  point  will  be  more  minutely  considered.  ^ ^ 

the  PRINCIPLES  OF  CRYSTALLOGRAPHY.  yv,  ^ 'b  t 

Solid,  homogeneous  bodies  are  composed  of  the  smallest  particles,  mole- 
cules, which  we  imagine  as  being  irregularly  placed  or  arranged  in  a 
regular  net-like  manner.  In  the  first  instance  they  are  amorphous,  show- 
ing like  physical  properties  in  every  direction,  while  in  the  second  case 
they  are  crystallhed,  manifesting  a similar  physical  deyiortment  in  all 
parallel  directions,  but  which  in  general  varies  in  different  directions. 


32 


inor(;anic  chemistry. 


In  crystallized  bodies  the  molecules  are  so  arranged  in  layers,  that  all 
the  particles  are  similar.  These  layers  form  i)lanes,  crystal  faces,  whose 
form  depends  on  the  nature  of  the  molecules;  in  crystals  differing  from 
one  another  chemically,  the  form  of  the  faces  is  also  different. 

A combination  of  all  the  crystal  faces  (planes)  circumscribing  a crys- 
tallized body  constitutes  its  crystal  form,  which  is  always  definite  for 
every  crystalline  chemical  comiiound.  I'heform  and  the  extension  of  the 
individual  faces  may  vary,  but  the  angles  produced  by  the  faces  remain 
unchanged.  (Law  of  the  constancy  of  angles  formed  by  crystal  faces.) 

A zone  is  produced  by  three  or  more  iilanes  cutting  one  another  in  paral- 
lel edges  ; the  zone  axis  is  the  direction  with  whi(  h the  edges  run  parallel. 

We  can  imagine  most  crystals  so  divided  by  planes  that  each  crystal 
face  upon  the  one  side  corres[)onds  to  a similar  face  on  the  ojjposite  side, 
producing  a like  angle  with  the  intersecting  plane.  Such  a i)lane  is  desig- 
nated di  symmetry-plane ; its  i)erpendicular  is  known  as  the  sy^nmetry  axis. 

The  symmetry  ratios  vary  with  the  individual  crystals.  One  may 
exhibit  greater  symmetry — more  symmetry-planes — than  another.  All 
those  crystalline  forms  in  which  there  is  a like  number  of  symmetry- 
planes  constitute  a system  of  crystallization. 

There  are  two  kinds  of  symmetry-planes.  Thus  crystals  exhibit  in 
part  one  or  several  symmetry-planes,  with  which  several  symmetry  axes 
are  parallel  and  which  may  be  exchanged  one  for  the  other  without 
altering  the  crystalline  form.  Symmetry  axes  of  this  description  are  said 
to  be  equivaleiit,  and  the  symmetry-planes  with  which  they  are  parallel 
are  called  the  principal  symmetry-pla?ies  and  their  perpendiculars  are  the 
principal  axes. 

Six  systems  of  crystallization  are  distinguished  on  the  basis  of  relations 
in  symmetry. 

I.  The  regular  or  isometric  system  with  three  principal  symmetry- 
planes  and  six  secondary  symmetry-planes. 

II.  The  hexagonal  system  with  one  principal  symmetry-plane  and  six 
symmetry-planes. 

III.  The  quadratic  or  tetrag07ial  systetn  with  one  principal  symmetry- 
plane  and  four  secondary  symmetry-planes. 

IV.  The  rhombic  system  with  three  secondary  symmetry-planes. 

V.  The  mo7ioclinic  system  with  one  secondary  symmetry-plane. 

VI.  The  triclinic  system,  in  which  there  is  no  symmetry-plane. 

I.  The  Isometric  System. — The  forms  of  this  system  are  referred 
to  an  axis  system  consisting  of  three  axes  of  equal  length  (^principal 
sym?fietry  axes)  and  at  right  angles  to  one  another.  Two  of  the  axes  lie 
in  a ])rincii)al  symmetry-})lane.  The  various  fundamental  forms  are  de- 
rived by  imagining  these  axes  a,  a,  a,  cut  by  planes  at  equal  or  unequal 
distances  from  tlieir  point  of  intersection.  The  axial  section  of  a plane 
is  termed  its  |)arameter.  If  the  smallest  of  these  be  designated  by  a,  and 
the  other  two  by  na  and  ma,  then  the  values  n and  m (the  coefficients  of 
the  parameters)  become  in  accordance  with  experience  rational  numbers. 
'I'here  are  seven  fundamental  forms: 

I.  Octahedron  {O)  (h'ig.  i).  'Fhe  faces  intersect  the  axes  at  equal 
distances  a : a : a from  the  center. 


INTRODUCTION. 


33 


2.  Cube  (ood?oo)  (Fig.  2).  The  faces  intersect  one  axis  and  are 

parallel  to  the  other  two  : a : : oo^r. 

3.  Dodecahedron  ( 00  (Fig.  3).  The  faces  intersect  two  axes  at  the  same 
distance  from  the  center,  while  they  are  parallel  to  the  third  axis  : a \ a:  ooa. 


4.  Trisoctahedron  (mO)  (Fig.  4).  The  faces  intersect  two  of  the  axes 
at  unity  and  the  third  at  a greater  distance  : a : a : ma. 

5.  Trapezohedron  or  Jcositetrahedron  {mOin)  (Fig.  5).  The  faces 
intersect  one  axis  at  unit  distance  and  the  other  two  at  equal  but  greater 
distances:  a : nia  : ma. 


34 


INORGANIC  CHEMISTRY. 


6.  Tetrahexahedro7i  ( o/^Oti)  (Kig.  6).  Tlic  faces  intersect  two  axes  at 
different  distances  and  the  thiixl  axis  at  infinity  : a : 7ux  : cc/'/. 

7.  Ilexoctaliedron  (xnOn)  (h'ig.  7).  The  faces  cut  all  three  axes  at 
different  distances  : a : ;ia  : xna. 

These  siixiple  forms  usually  occur  in  combination;  this  is  also  true  of 
the  remaining  systems.  Thus  Fig.  8 rej)resents  the  combination  of 
dodecahedron  ( 006?),  cube  ( oo(9co),  and  octahedron  (6^). 

II.  The  Hexagonal  System. — 'I'here  are  seven  symmetry-j)lanes  in 
this  system,  six  of  which  are  at  right  angles  to  the  seventh — the  iirinciiial 
symmetry  jilane.  The  six  planes  referred  to  cut  one  another  at  an  angle 
of  30°.  The  intersecting  lines  of  three  alternating  ])lanes  (cutting  one 
another  at  60°),  together  with  the  princii)al  symmetry-j)lane,  are  regarded 
as  the  secondary  axes  (these  are  equal  a : a : a).  The  fourth  axis  is  the 
axis  (c)  at  right  angles  to  the  principal  symmetry-plane.  The  ratio  : c 
is  irrational  and  is  very  definite  for  every  substance  crystallizing  in  the 
hexagonal  system. 


P 


Fig.  9. 


The  forms  of  this  system  are  either  double  consisting  of  twelve 

or  six  faces  converging  above  and  below  at  the  terminals  of  the  principal 
axis,  or  they  are  twelve-  or  six-sided  prisms,  the  faces  of  which  are  parallel 
to  the  principal  axis.  To  these  must  be  added  the  pair  of  faces  perpen- 
dicular to  the  principal  axis — the  basal  planes.  Fig.  9 represents  the 
hexagonal  pyra7nid  (^P)  and  Fig.  10  the  hexagonal pyra77iid  (^P)  with  the 
pris77i  ( ooP).  The  angle  produced  by  the  intersection  of  two  prism  faces 
ecpials  60°. 

III.  Tetragonal  or  Quadratic  System. — This  system  like  the 
hexagonal  is  characterized  by  a single  principal  symmetry-plane  to  which 
four  alternating  similar  symmetry-planes  cutting  one  another  at  45°  are 
fierpendicular.  d'here  are  three  axes  at  right  angles  to  one  another. 
'I'he  ))rinci])al  axis  (c)  is  normal  to  the  principal  symmetry-plane,  while 
the  other  two  (a,  a)  correspond  to  the  intersecting  directions  made  by  two 
ecpial  symmetry-])]ancs  with  the  principal  symmetry-plane. 

'The  forms  of  this  system,  as  in  the  hexagonal  system,  consist  of  pyra- 
mids, jirisms,  and  basal  i)lane  (})inacoid).  The  pyramids  are  made  up  of 


INTRODUCTION. 


35 


four  or  eight  faces  above  and  below.  The  i)risms  also  have  four  or  eight 
laces;  e.  Fig.  ii:  tetragonal  pyratnid  (^P)\  Fig.  12:  tetragonal 
prism  ( 00^)  with  the  pyramid  {P").  The  angle  formed  by  two  prism 
faces  is  90°. 

IV.  The  Rhombic  System. — In  this  system  there  are  three  dis- 
similar symmetry-planes  at  right  angles  to  one  another.  Their  directions  of 
intersection  are  assumed  to  be  the  axes  a,  I?,  c.  The  rhombic  pyramid  (P) 
(Fig.  13)  is  the  fundamental  form.  It  cuts  the  three  axes  (of  unequal 
lengths)  at  their  unit  distances  (^zis  the  brachy-axis,  b the  macro-axis,  and 
c the  vertical  axis').  There  are  other  pyramids  in  which  the  faces  intersect 


Fig.  12. 


Fig.  13. 


Fig.  14. 


the  axes  at  other  than  unit  distances  from  the  center.  Prismatic  forms  re- 
sult when  the  faces  are  parallel  to  one  axis.  If  they  are  parallel  to  c the 
rhombic  prism  ( ccP)  (Fig.  14),  and  macro-  or  brachydomes  if  the  faces 
run  parallel  to  b or  a,  e.  g.,  Fig.  15,  the  7nacrodome. 

The  pair  of  faces  perpendicular  to  the  vertical  axis  constitutes  the  basal 
planes,  while  the  other  pairs  of  faces  at  right  angles  to  the  other  two  axes 
are  termed  the  macro-  and  brachy-pinacoids. 

V.  The  Monoclinic  System. — There  is  but  one  symmetry-plane  in 
this  system.  It  is  directed  toward  the  observer  when  he  examines  the 


36 


INORGANIC  CHEMISTRY. 


crystals.  The  axis  system  consists  not  (mly  of  tlie  perpendicular  (/^-axis) 
to  the  syminetry-])lane,  but  also  (d' the  intersecting  lines  und  c),  \vhi(  li 
form  two  crystal  faces  with  and  at  right  angles  to  the  symmetry-plane  and 
are  inclined  to  one  another  at  any  arbitrary  angle  (fi).  The  vertical  axis 
is  c,  while  a,  directed  toward  the  observer,  is  designated  the  clino-axis. 
The  forms  are  like  those  of  the  rhombic  system,  l^iach  ])yramid,  Innvever, 
resolves  itselt  into  two  independent / these  are  distinguished 
as  positive  (-(-  P)  and  negative  ( — P)  (Fig.  i6).  d'here  are  positive  and 
negative  orthodomes,  each  arising  from  a jiair  of  faces,  which  correspond 
to  the  rhombic  macrodomes.  The  prismatic  form,  i)arallel  to  the  clino- 
axis,  is  termed  the  cli nodome. 

VI.  Triclinic  System. — Any  three  crystal  planes  are  selected  for 
the  axis  planes.  Tneir  intersecting  lines  constitute  the  axes  zz  (brachy- 
diagonal),  b (macrodiagonal),  c (vertical  axis).  All  the  axes  intersect 
one  another  at  oblicpie  angles.  The  angle  between  b and  c is  designated 
a,  that  between  a and  c (3,  and  that  between  a and  b y.  There  is  no 
symmetry-plane  present;  each  parallel  pair  of  faces  constitutes  a crystal 
form,  so  that  a pyramid  intersecting  the  three  axes  at  the  distances  a:  b:  c 
consists  of  four  different  crystal  forms  which  are  marked  as  in  Fig.  17 
with  P',  P^,  'P,  ^P.  All  the  other  forms,  as  in  the  monoclinic  system,  are 
derived  from  these,  with  the  difference  that  in  this  system  the  distinction 
between  right  and  left  must  be  observed. 

In  nature,  crystals  rarely  occur  so  regularly  developed  as  represented 
in  the  preceding  forms.  Usually  they  are  more  or  less  elongated  in  one 
or  more  directions  and  this  causes  the  faces  of  one  and  the  same  form  to 
become,  as  regards  their  spacial  extension,  unequal  and  the  whole  crystal 
appears  to  be  distorted.  Such  a distortion,  however,  never  influences 
the  position  of  the  faces  with  reference  to  .the  axes.  The  angles  produced 
by  the  faces  remain  unchanged  so  long  as  the  temperature  is  the  same. 
(Law  of  the  constancy  of  the  angles  formed  by  the  faces.)  The  measure- 
ment of  the  angle  formed  by  the  faces  by  means  of  the  goniometer  is  the 
only  means  we  possess  of  unraveling  complicated  combinations,  and  even 
this  aid  is  not  in  all  cases  satisfactory  in  determining  definitely  and  surely 
that  a form  belongs  to  one  or  another  system.  For  this  purpose  a careful 
investigation  of  the  physical  properties  is  frequently  essential,  because  it 
occurs  that  crystals  apparently  develop  a higher  symmetry  geometrically 
than  belongs  to  them  when  their  physical  properties  are  considered 
(pseudo-symmetric  crystals). 

Another  irregularity  in  their  development  consists  in  the  fact  that  the 
faces  themselves  are  not  even  and  smooth.  They  frequently  occur  bent, 
e.  g.,  in  the  diamond  and  dolomite;  drusy,  due  to  the  protuberance  of 
numerous  little  solid  angles,  uj)on  the  faces,  of  other  forms  differently 
orientated,  or  striated  by  the  frequent  alternating  appearance  of  two 
faces— g.,  the  striation  of  the  cube  faces  of  pyrite  due  to  the  oscillatory 
a])))earancc  of  a ])cntagonal  dodecahedron — rough,  corroded,  etc.,  etc. 

'I’wo  or  more  crystals  are  often  grown  together;  this  growth  being 
either  regular  or  irregular.  When  regular  the  crystals  are  ]xarallel  to  one 
aiu>ther  and  are  designated  pa7’anel  growths,  or  they  are  not  ])arallel  but 
yet  limited  by  a very  definite  regularity.  Such  growths  are  called  twins. 


INTRODUCTION. 


37 


Their  regularity  consists  in  the  fact  that  the  two  crystals  forming  the 
twin  are  symmetrically  developed  with  one  another  in  relation  to  a crys- 
tallographically  possible  plane — the  iw inning-plane.  Fig.  i8  represents 
an  example  of  such  a twin  crystal.  Two  octahedra,  or  better,  two  half 
octahedra,  are  here  symmetrical  in  reference  to  an  octahedral  face. 
Spinel  and  magnetite  are  examples  of  this  class.  The  plane  along  which 
the  two  individuals  have  developed  is  the  growth-plane.  It  is  not  neces- 
sary that  this  should  at  the  same  time  be  the  twinning-plane,  for  two 
crystals  can  be  symmetrical  to  one  another  in  referenc  e to  one  plane  and 
yet  be  developed  along  another  plane.  This,  for  example,  is  often  the  case 
with  what  are  termed  the  Carlsbad  orthoclase  twins.  In  them  the  clino 
pinacoid  is  the  growth-plane  and  the  orthopinacoid  the  twinning-plane. 


'T'/ 

Fig.  I 8. 


Fig.  19. 


oP 

Fig.  21. 


Fig.  20. 


Crystals  may  also  grow  on  both  sides  beyond  the  growth-plane, 
tration-twins  or  crosses  are  the  result. 

In  addition  to  the  simple  forms  appearing  with  all  their  faces,  hence 
termed  holohedral  for7n^  others  occur  in  which  only  half  the  possible 
faces  are  present — the  heniihedral  form.  These  result  if  we  imagine  a 
holohedral  form  to  be  divided  by  symmetry-planes  into  congruent  parts, 
and  then  have  one  of  every  two  parts,  which  are  symmetrical  in  reference 
to  these  planes,  fall  away.  In  this  manner  every  holohedral  figure  will 
yield  two  heniihedral  forms  which  differ  from  one  another  in  their 
position,  and  are  either  congruent  or  symmetrical.  Thus,  the  octahedron 
of  the  isomeric  system  yields  the  teti'aliedron,  Fig.  19,  and  the  hexagonal 
pyramid,  the  rhombohedron  (Fig.  20). 

A comparatively  rare  phenomenon  has  been  observed  in  the  crystals  of 


38 


INORGANIC  CHEMISTRY. 


some  substances — namely,  a different  development  at  the  opposite  ends  of 
an  axis.  This  is  known  as  he77iimorphi5m.  Struvite  (magnesium  ammo- 
nium phosphate  Mg(NH4)r04  + 6 HjO),  Fig.  21,  is  an  exampleof  this. 
On  the  upper  end  of  the  vertical  axis  occur  the  domes  P 00  and  P ^ , 
while  below  there  is  only  the  dome  P^  and  the  basal  plane  oP.  Tour- 
maline, calamine,  cane  sugar,  etc.,  exhibit  the  same  phenomenon. 

Substances  crystallizing  variously  in  the  same,  or  in  two  or  three 
different  systems  are  said  to  be  di77iorphous,  tnTtiorphous,  etc.  Titanium 
dioxide  (TiOa),  for  example,  is  found  in  nature  in  crystals  of  the  quadratic 
system  as  the  minerals  anastase  and  rutile,  and  in  those  of  the  rhombic 
system  as  brookite  (see  also  Sulphur). 

Various  substances  which  crystallize  in  similar  or  in  forms  very  much 
alike  in  the  same  system  are  called  iso77io7phous  {sqq  also  Isomorphism). 


SPECIAL  PART. 


■ - " CLASSIFICATION  OF  THE  ELEMENTS. 

Ordinarily  we  are  accustomed  to  divide  the  elements  into  two  groups  : 
metals  and  non-metals  (seep.  19).  The  former  possess  metallic  appear- 
ance, are  good  conductors  of  heat  and  electricity;  the  latter,  the  metal- 
loids or  non-metals,  do  not  have  these  properties,  or  at  least  in  less  degree. 
In  chemical  respects  the  metalloids  have  the  tendency  to  combine  with 
hydrogen,  forming  volatile,  generally  gaseous,  compounds;  their  oxygen 
derivatives  form  acids  with  water.  The  metals,  on  the  contrary,  do  not 
unite  with  hydrogen,  or  at  least  do  not  form  volatile  compounds  with  it, 
and  their  oxygen  derivatives  yield  chiefly  the  so-called  bases  with  water. 
Further,  the  compounds  of  metals  with  the  non-metals  are  so  decomposed 
by  the  electric  current  that  the  metal  separates  at  the  electro-negative, 
and  the  non-metal  at  the  electro-positive  pole.  From  this  we  observe 
the  metals  are  more  electro-positive — more  basic;  the  metalloids  more 
electro-negative — of  an  acid-forming  nature.  A sharp  line  of  difference 
between  metals  and  metalloids  does  not  exist.  There  are  elements,  like 
antimony,  which  in  their  external  appearance  resemble  metals,  while  in 
a chemical  respect  they  deport  themselves  throughout  as  metalloids,  and 
vice  versa.  Thus  hydrogen,  a gaseous  element,  is  like  the  metals  in  its 
entire  chemical  character,  while  metallic  antimony  arranges  itself  with 
the  metalloids. 

It  is  therefore  better  to  divide  the  elements  into  separate  natural  groups, 
based  upon  their  chemical  analogies.  The  best  and  only  correct  classifi- 
cation of  all  the  elements  depends  on  the  law  of  periodicity,  according 
to  which  the  properties  of  the  elements  and  of  their  compounds  present 
themselves  as  a periodic  function  of  the  atomic  weights.  Later  we  shall 
treat  of  the  periodic  system  more  at  length ; it  forms  the  basis  of  this 
text-book,  and  in  accordance  with  this  doctrine  we  consider  the  elements 
in  single  natural  groups  of  similar  chemical  deportment.  The  first  of 
these  groups,  comprising  almost  all  the  so-called  non  metals,  are  the 


following : 

Fluorine 

Oxygen 

Nitrogen 

Carbon 

Boron 

Chlorine 

Sulphur 

Phosphorus 

Silicon 

Bromine 

Iodine 

Selenium 

Tellurium 

Arsenic 

Antimony 

. Hydrogen  does  not  belong  to  any  of  these  groups;  uniting  the  metal- 
lic and  non-metallic  characters  in  itself,  it  represents,  as  it  were,  the  type 
of  all  elements,  and  therefore  it  will  receive  first  attention.  Boron 

39 


40 


INORGANIC  CHEMISTRY. 


occupies  an  isolated  position.  It  has  been  classed  with  the  non-metals, 
but  differs  somewhat  from  them  in  chemical  deportment.  It  forms  the 
transition  to  the  metallic  elements,  beryllium  and  aluminium. 


HYDROGEN. 

Atom:  H = i.oi.  Molecule:  Hg  = 2.02. 

Hydrogen  (hydrogeniuin),  a gaseous  body,  occurs  rarely  in  a free  con- 
dition upon  the  earth’s  surface, — in  the  gases  from  volcanoes  and  in 

those  issuing  from  the  earth, 
— as  an  enclosure  in  minerals, 
a product  of  decay,  and,  from 
recent  statements,  it  is  found 
in  very  small  amount  in  the 
atmosphere.  It  is,  however, 
present  in  considerable  quan- 
tity in  the  photosphere  of  the 
sun  and  fixed  stars.  In  com- 
bination, it  is  found  chiefly 
as  water,  and  in  substances  of 
vegetable  and  animal  origin. 
Paracelsus  first  observed  this 
element  in  the  sixteenth  cen- 
tury, and  called  it  inflamma- 
ble air;  and  in  1766  Cav- 
endish, recognizing  it  as  a 
peculiar  gas,  named  it  inflam- 
mable air.  In  1 783  Lavoisier 
proved  that  hydrogen  was  a 
constituent  ot  water — a chemical  compound  of  the  elements  hydrogen 
and  oxygen — by  conducting  steam  over  ignited  metallic  iron. 

Preparation. — It  may  be  readily  obtained  from  water.  The  decompo- 
sition of  the  same  by  the  removal  of  oxygen  can  be  effected  by  some  metals, 
like  sodium  and  potassium,  at  the  ordinary  temperature.  Both  metals  act 
very  energetically  upon  it,  liberating  gaseous  hydrogen.  To  perform  the 
experiment,  take  a piece  of  sodium,  roll  it  up  in  a piece  of  wire  gauze, 
and  shove  it,  with  nippers,  under  the  mouth  of  a glass  cylinder  filled  with 
and  inverted  over  water  (Fig.  22).  Bubbles  of  hydrogen  are  at  once 
disengaged,  displace  the  water  and  collect  in  the  upper  part  of  the  cylin- 
der. The  reaction  occurring  between  the  sodium  and  water  is  expressed 
by  the  following  chemical  equation  : 

11,0  -f  Na  NaOH  + H. 

Water.  Sodium.  Hydrogen, 

I he  compound  NaOH,  known  as  sodium  hydroxide,  remains  dissolved  in 
the  excess  of  water. 

Other  metals  decom))ose  water  in  a similar  manner,  at  an  elevated 
temperature.  To  effect  this  with  iron  allow  steam  to  pass  through  a tube 


CLASSIFICATION  OF  THE  ELEMENTS.  41 

filled  with  iron  filings,  exposed  to  a red  heat  in  a combustion  furnace, 
d'he  iron  withdraws  oxygen  from  the  water,  combining  with  it,  while  the 
hydrogen  set  tree  is  collected.  Magnesium  powder  reacts  similarly,  but 
at  a much  lower  temperature,  upon  steam  (Ber.  26  (1893),  I,  59). 

For  laboratory  purposes,  hydrogen  is  prepared  by  the  action  of  zinc 
upon  hydrochloric  or  sulphuric  acid.  The  reaction  with  the  latter  acid 
is  as  follows : 

Zn  + H^SO,  ZnSO,  + H.,. 

Sulphuric  acid.  Zinc  sulphate. 

Place  granulated  zinc  (obtained  by  dropping  molten  zinc  into  water) 
in  a double-necked  flask  (Fig.  23),  and  introduce  sulphuric  acid  (diluted 
with  about  3 vols.  of  H2O)  through  the  funnel  tube,  b.  The  liberation 
of  gas  begins  immediately,  and  the  hydrogen,  escaping  through  the  exit 
tube,  f,  is  collected  as  previously  described. 


Fig.  23. 

The  hydrogen  thus  formed  has  a faint  odor  due  to  a slight  admixture 
of  foreign  substances  (the  hydrides  of  sulphur,  arsenic,  phosphorus 
and  carbon — if  these  elements  are  contained  in  the  metal  used  in  the  gas 
evolution).  It  is  therefore  conducted  through  a solution  of  potassium 
permanganate  to  purify  it. 

Many  other  metals,  e.  g.,  iron,  behave  like  zinc  with  dilute  acids. 
Some  metals — zinc,  iron,  and  aluminium — dissolve,  when  finely  divided, 
in  sodium  or  potassium  hydroxide  with  the  liberation  of  hydrogen  : 

Zn  -f  2NaOH  = Zn(ONa)2  + 

Pure  hydrogen  may  be  obtained  by  heating  potassium  formate  with 
potassium  hydroxide:  CHO^K  -f-  KOH  =:  K2CO3  -(-  H2 ; for  technical 
purposes  by  heating  zinc  or  iron  with  calcium  hydroxide  (slaked  lime)  in 
a combustion  tube  : Zn  -(-  Ca02H2  = ZnO  -j-  CaO  -f-  Hj,  or  with  coal 
(anthracite)  : 

2Ca02ll2  + c = CaCOg  + CaO  -f-  2H2. 

4 


42 


INORGANIC  CHEMISTRY. 


A very  important  method  for  tlie  i)rei)aration  of  hydrogen  consists  in 
decomposing  certain  arpieons  solutions  by  means  (jf  tlie  electric  current. 
This  decom])osition  usually  proceeds  as  if  the  water  liroke  down  into  its 
constituents  : 2H./ ) — 2I I2  -|-  0.2.  The  hydrogen  ajipears  at  the  negative 
pole.  The  electrolysis  of  such  solutions  will  be  discussed  exhaustively  later. 


Purifying  atuf  D)yjng  of  Gases. — To  free  gases  of  the  substances  mechanically  carried 
along  during  their  disengagement,  it  is  best  to  conduct  them  through  variously  constructed 
wash-bottles,  filled  with  water  or  liipiids,  which  will  absorb  the  impurities.  Ordinarily  the 
so-called  WoullT  bottles  are  employed  (compare  Figs.  31  and  35).  'J’he  open  tube, 

placed  in  the  middle  tubulure,  is  called  the 
safety  tube.  It  serves  to  equalize  the  inner 
pressure  with  that  of  the  external  atmos- 
phere. Wash-bottles  of  various  construction 
will  be  represented  in  the  several  sketches. 

Gases  liberated  from  an  aqueous  liquid 
are  always  moist,  as  they  contain  aqueous 
vapor.  To  remove  this  conduct  them 
through  vessels  or  tubes  filled  with  hygro- 
scopic substances  (see  Fig.  30).  Calcium 
chloride,  burnt  lime,  sulphuric  acid,  etc., 
are  used  for  this  purjjose. 

Apparatus  for  the  Gene?-ation  and  Col- 
lection of  Gases. — In  the  ap])aratus  pictured 
in  Fig.  23,  the  liberation  of  hydrogen  con- 
tinues uninterruptedly  as  long  as  zinc  and 
sulphuric  acid  are  present.  To  control  the 
generation  of  the  gas  we  have  recourse  to 
different  forms  of  apparatus.  One  of  the 
most  practicable  of  these  is  that  of  Kipp. 
It  consists  of  two  glass  spheres,  d and 
Fig.  24,  in  the  upper  opening  of  w'hich 
there  is  a third  sphere,  c,  fitting  air-tight 
and  provided  with  an  elongated  tube. 
It  serves  as  a funnel.  Granulated  zinc  is 
placed  in  the  middle  sphere  through  the 
tubulure  <?,  and  dilute  sulphuric  acid  is 
poured  into  the  spherical  funnel,  which 
first  fills  d,  then  ascends  to  where  it  comes  in  contact  with  the  zinc ; at  once  the  evo- 
lution of  hydrogen  commences  and  the  gas  escapes  through  e.  Upon  closing  the  stop- 
cock of  the  tube  fixed  in  e,  the  hydrogen  which  is  set  free  presses  the  sulphuric  acid  out  of 
b,  and  consecpiently  the  liberation  of  the  gas  ceases.  On  again  opening  the  stop-cock,  the 
acid  rises  in  b to  the  zinc,  and  the  evolution  of  gas  commences  anew.  The  vessel  a con- 
tains water  or  concentrated  sulphuric  acid  to  wash  the  escaping  hydrogen. 

The  .somewhat  complicated  Kipp  apparatus  may  be  advantageously  replaced  by  the 
following  siinjde  contrivance  recommended  by  Debray  (Fig.  25),  Two  bottles  provided 
with  openings  near  their  bottom,  in  which  are  glass  tubes,  are  connected  by  a rubber  tube. 
4'he  bottle  A is  filled  with  granulated  zinc,  and  B with  dilute  sulphuric  acid.  The  cock 
R clo.ses  A.  When  this  is  opened  the  sulphuric  acid  flows  from  B to  A,  to  the  zinc, 
and  the  evolution  of  gas  commences.  D represents  a form  of  wash-bottle  suggested 
by  Kemi)f.  The  gas  evolution  can  be  quickly  regulated  by  raising  or  lowering  the 
Inktle  /)’. 

'fhe  Mohr  ap])aratus  ( Fig.  26)  is  very  convenient  for  the  evolution  of  gases.  Into  the 
second  cylinder  containing  the  acid  another  narrower  cylinder  closed  by  a stop-cock  is 
introduced,  'fhe  zinc  (iron,  marble,  etc.)  is  ]daced  above  the  contraction  of  the  inner 
cylinder.  On  opening  the  stop-cock  the  acid  enters  aiuk evolution  of  gas  commences. 
On  closing  the  stopcock  the  hydrogen  pres.ses  the  acid  back  from  the  inner  cylinder;  tlae 
evolution  of  gas  cease.s. 


S 


Fig.  24. 


CLASSIFICATION  OF  'I  HE  ELEMENTS. 


43 


Gasometers  of  various  construction  serve  to  collect  and  preserve  gases.  In  Fig.  27  we 
have  the  ordinary  gasometer  of  Pepys.  It  is  constructed  of  sheet  copper  or  zinc,  and 
consists  of  two  cylindrical  vessels,  the  lower  one  closed,  the  upper  open,  communicating 
with  each  other  by  the  two  tubes  a and  d.  The  tube  c is  only  a support.  To  this  end, 
pour  water  into  the  upper  cylinder,  and  open  a and  tr  ; the  water  then  flows  through  a, 
nearly  reaching  the  bottom  of  the  lower  cylinder,  while  the  air  escapes  through  e.  When 
the  lower  cylinder  is  filled  with  water  close  a and  e (the  last  traces  of  air  can  be  removed 
by  opening  b').  To  fill  the  gasometer  with  gas,  remove  the  cover  of  the  side  tubulure 
d,  and  introduce  the  tube  from  which  the  gas  is  escaping.  The  latter  rushes  up  into  the 
cylinder,  while  the  water  flows  out  the  tubulure.  When  the  water  is  displaced  by  the 
gas,  close  d,  after  filling  the  upper  cylinder,  and  then,  if  desired,  open  a,  and  the 
gas  can  be  set  free,  either  by  e or  b. 

In  addition  to  the  gasometer  described,  various  other  forms  are  employed  ; gas-bags 
and  tubes  are  very  well  adapted  for  preserving  gases. 

Physical  Properties. — Hydrogen  is  a colorless,  odorless,  and  tasteless 
gas.  It  conducts  heat  and  electricity  better  than  all  other  gases.  This 


Fig.  25. 


Fig.  26. 


may  be  readily  proved  by  the  following  experiment : A current  of  elec- 
tricity is  sent  through  a thin  platinum  spiral,  and  while  the  latter  remains 
in  the  air  or  some  other  gas,  it  will  glow,  but  in  an  atmosphere  of  hydro- 
gen the  spiral  will  not  become  luminous,  or  at  once  cease  glowing. 
The  metallic  nature  of  hydrogen  is  shown  by  this  behavior. 

Of  all  gases  hydrogen  is  the  most  difficult  to  liquefy,  because  its  criti- 
cal temperature  is  the  lowest  (about  — 220°);  it  must,  therefore,  be  ex- 
posed to  the  most  intense  cold. 

The  procedure  first  adopted  by  Olszewski  and  more  recently  by  Dewar 
for  the  production  of  liquid  hydrogen  (also  ajiplied  to  other  gases)  con- 
sists in  suddenly  lowering  the  pressure  of  the  sufficiently  cooled  and 
powerfully  compressed  gas,  when  it  will  rapidly  expand  and  a portion  of 
it  will  deprive  another  portion  of  so  much  heat  that  the  second  portion 
will  liquefy.  Dewar  permitted  hydrogen  cooled  to  — 205°  and  under  a 
pressure  of  180  atmospheres  to  enter  a vacuum-vessel  also  cooled  to  — 200°. 


44 


INORGANIC  CHKMISTKY. 


'File  liquid  hydrogen  i)rei)ared  in  this  way  is  colorless,  very  nioliile  and 
has  the  specific  gravity  of  0.07.  It  boils  at  about  —240°  under  the  ordi- 
nary atmospheric  pressure. 

Like  all  gases,  coercible  with  difficulty,  hydrogen  is  but  slightly  soluble 
in  water,  1000  volumes  dissolving  19  volumes  of  the  gas  at  14°  C.  'Fhe 
coefficient  of  absori)tion  of  hydrogen  by  water  is  therefore  0.0193.  It  is 
the  lightest  of  all  gases.  That  is,  a definite  volume  of  it  weighs  less  than 
an  equal  volume  of  any  other  gas  at  the  same  temperature,  under 

like  iiressure,  and  subject  to  a like 
action  of  gravity.  It  is  generally 
understood  that  the  comparison 
of  the  weight  of  equal  gas  volumes 
is  made  at  0°,  760  mm.  barometric 
pressure  at  sea-level  and  in  the  45° 
of  geographical  latitude.  These 
are  normal  conditions,  and  under 
them  I liter  of  hydrogen  weighs 
0.08988  gram,  i liter  of  oxygen: 
1. 4291  grams,  and  i liter  of  air: 
1.2930  grams.  The  thousandth 
part  of  each  number,  the  weight  of 
one  cubic  centimeter,  is  the 
specific  gravity  of  the  gas  referred 
to  water  at  4°.* 


Fig.  27.  Fig.  28. 

From  this  it  follows  that  air  is  14.4  (=  times  heavier  than 

hydrogen,  and  oxygen  15.9  ^ ) times  heavier  than  the  latter. 

'Fhese  numbers  are  the  densities  of  these  gases  referred  to  hydrogen  as  unit. 
If,  however,  the  Aveight  of  the  air  volume  is  selected  as  unit  we  obtain 
0.06951  (—  ^1^.2  9 3^)  density  of  hydrogen,  and  1.T052  (— 

as  that  of  oxygen.  It  has  become  the  rule  quite  recently  in  chemical 
discussions  to  refer  densities  to  that  element  which  is  taken  as  the  standard 
in  atomic  weights  : oxygen,  and  for  reasons  which  will  be  given  subse- 


* f'or  llic;  (let(‘nnitiatif)ii  of  these  numbers  compare  Morley,  Z.  j^hys.  C^i.  17  (1895') 
87;  20  (]H()G)  68,  242,  417,  and  4'Iiom.sen,  Z.  f.  anorg.  Cli.  12  (1896)  I ; also  Kohl- 
rau.sch,  Lcilfaden  der  prakt.  Physik,  8 Aull.  (1896),  459.  Morley  gives  as  the  most  prob- 
able values  for  the  weight  of  a liter  of  liydrogeii  and  that  of  oxygen  : 0.089873  and 
1.4290;  'riuun.sen  : 0.089947  and  1.42906. 


CLASSIFICATION  OF  THE  EI.EMENTS.  45 

quently,  the  weight  of  its  volume  is  placed  at  32.  Then  for  hydrogen 
there  would  result  the  number  2.0126  and  for  air  28.954. 

The  volume  expressed  in  cubic  centimeters  which  i gram  of  gas  would 
occupy  under  normal  conditions  is  its  specific  volume.  If  the  weight  of  one 
cubic  centimeter  of  gas  is  represented  by  a,  its  specific  volume  would  be 
Vgrrr-T-  Heuce  the  specific  volume  of  hydrogen  is  = 1 1, 1 26  ; 

that  is,  one  gram  of  hydrogen  under  normal  conditions  occupies  11,126 
cubic  centimeters  ( 1 1. 1 2 liters).  The  specific  volume  of  air  is  q ^ 293 
= 773-4  c.c,,  and  that  of  oxygen  =699-7  c.c. 

That  hydrogen  is  lighter  than  air  is  shown  by  a balloon  of  collodion 
or  gum  filled  with  the  former  rising  in  the  latter ; this  can  also  be  seen 
in  soap  bubbles  filled  with  hydrogen.  In  consequence  of  its  levity, 
hydrogen  may  be  collected  in  inverted  vessels  (opening  turned  down)  by 
displacing  the  air,  and  can  also  be  poured  from  one  cylinder  into  another, 
as  represented  in  Fig.  28.  The  hydrogen  flows  from  the  inclined 
cylinder  into  the  one  held  vertically  and  filled  with  air,  which  it  expels. 
Such  a separation  of  gases,  based  on  their  varying  specific  gravity,  is  only 
temporary,  as  they  soon  mingle  with  each  other  by  diffusion.  By  virtue 
of  its  levity  and  mobility,  which  the  kinetic  gas  theory  attributes  to  the 
great  velocity  of  the  gas  ])articles,  hydrogen  penetrates  porous  bodies  with 
ease,  and  diffuses  througli  both  animal  and  vegetable  membranes,  as  well 
as  through  gutta-percha  (consult  Air,  upon  diffusion  of  gases).  Metals, 
e.  g.,  iron,  platinum,  palladium,  permit  a free  passage  to  hydrogen,  when 
they  are  raised  to  a red  heat.  They  are  impenetrable  to  other  gases  (see 
Dissociation  of  Water).  This  behavior  is  in  part  probably  dependent  upon 
the  chemicalattraction  of  these  metals  for  hydrogen. 

Chemical  Properties. — Hydrogen  is  character- 
ized by  its  ability  to  burn  in  the  air,  when  it  com- 
bines with  the  oxygen  of  the  latter  and  forms 
water;  hence  its  name  hydrogenium,  due  to  Lavo- 
isier (from  water,  and  y&wdoi,  I produce). 

Its  flame  is  faint  blue,  and  almost  non-luminous, 
but  possesses  a very  high  temperature.  When  a 
mixture  of  hydrogen  and  air  is  ignited  a violent 
explosion  ensues ; therefore,  before  bringing  a 
light  in  the  vicinity  of  hydrogen  disengaged  in  a 
vessel  filled  with  air,  allow  the  latter  to  escape 
completely,  otherwi.se  the  vessel  will  be  shattered 
to  pieces  by  the  explosion. 

As  hydrogen  itself  is  inflammable,  it  cannot 
sustain  the  combustion  of  other  bodies  which  will 
burn  in  the  air.  If  a burning  candle  be  introduced  into  an  inverted 
cylinder  containing  the  gas  (Fig.  29)  the  latter  will  ignite  at  the  mouth 
of  the  vessel,  but  the  candle  will  be  extinguished. 

Water  is  the  product  of  the  combustion  of  hydrogen  in  the  air.  It  is 
a chemical  compound  containing  hydrogen  and  oxygen.  To  render  the 
formation  of  it  visible,  by  the  combustion  of  hydrogen,  the  flame  of  the 
latter  is  made  to  burn  under  a cold  glass  jar  (Fig.  30).  The  sides  of  the 


46 


INORGANIC  CHKMISTRY. 


latter  are  soon  covered  vvitli  moisture,  whicli  collects  in  drojis.  d'o  avoid 
any  deception  the  hydrogen  is  first  conducted  through  sulphuric  acid  or  a 
tube  filled  with  calcium  chloride,  to  absorb  all  moisture. 

At  the  ordinary  temperature  hydrogen  and  oxygen  do  not  react  with- 
out the  intervention  of  a third  substance. 

The  union  of  hydrogen  with  oxygen  to  form  water  occurs  at  about 
200°  (compare  p.  28),  if  brought  in  contact  with  aflame,  or  by  the  pas- 
sage of  an  electric  sjiark  through  the  mixture.  The  combination  can  be 
effected  at  ordinary  temperatures  with  the  aid  of  platinum  sponge  or  ])al- 
ladium  (platinum  sponge,  palladium  asbestos)  ; the  jilatinum  sponge  con- 
sists of  finely  divided  metal,  obtained  by  the  ignition  of  ammonio- 
idatinum  chloride  (see  Platinum).  If  a stream  of  hydrogen  be  directed 
upon  a piece  of  freshly  ignited  platinum  sponge,  the  gas  will  at  once 
ignite.  This  is  due  to  the  power  of  the  platinum  to  condense  hydrogen 


Fig.  30, 


and  oxygen  upon  its  surface,  and  thereby  increase  their  ability  to  unite 
(^Dxbereiner  lamp). 

The  absorption  of  hydrogen  by  the  metal  palladium  is  very  character- 
istic. As  already  known,  water  is  so  decomposed  by  the  electric  current 
that  hydrogen  separates  at  the  electro-negative  pole  and  oxygen  at  the 
electro-positive.  Now,  if  a piece  of  palladium,  in  sheet  or  wire  form,  be 
attached  to  the  electro-negative  pole,  the  disengagement  of  hydrogen 
does  not  occur,  because  it  is  absorbed  by  the  palladium,  in  a quantity 
about  eight  hundred  and  fifty  times  the  volume  of  the  latter.  Palladium 
also  absorbs  hydrogen  when  it  is  heated  to  200°.  The  palladium  expands, 
becomes  lighter  in  weight,  but  retains  its  metallic  appearance.  Its 
tenacity  and  ])ower  of  conducting  heat  and  electricity  are  but  little  im- 
paired. 'The  compound  of  palladium  and  hydrogen,  therefore,  conducts 
itself  like  an  alloy  of  two  metals.  From  the  specific  gravity  of  the  com- 
f)ound  (according  to  Oraham),  the  specific  gravity  of  the  condensed 
hydrogen  is  found  to  be  0.62  (water  i),  and  is,  therefore,  somewhat 


CONDENSATION  OF  GASES. 


47 


heavier  than  the  metal  lithium  (compare  Palladium).  The  metals 
potassium  and  sodium  absorb  hydrogen  when  heated  from  200°  to  400°, 
forming  alloys  (Na4H2  and  K4H2)  in  which  the  density  of  hydrogen  is 
again  equal  to  0.62.  The  experiments  of  Winkler  and  of  Moissan  indi- 
cate that  Other  metals  also  unite  with  hydrogen:  barium,  calcium,  stron- 
tium, yttrium,  lanthanum,  cerium,  and  thorium.  These  facts  prove  the 
metallic  character  of  hydrogen,  which  is  also  indicated  by  its  ability  to 
conduct  heat  and  electricity.  Later,  we  will  observe  that  this  element 
displays  the  character  of  a metal  in  its  entire  chemical  deportment,  and 
that  it  must  be  regarded  as  a gaseous  metal  at  ordinary  temperatures. 


CONDENSATION  OF  GASES. 

CRITICAL  CONDITION.  0 

Hydrogen  and  several  other  gases  (oxygen,  nitrogen,  carbon  monoxide, 
methane,  nitric  oxide),  were  considered  until  1877  non-condensable — 
permanent  gases,  inasmuch  as  all  attempts  to  liquefy  the  same  were  fail- 
ures, notwithstanding  Natterer  (1852)  had  employed  a pressure  of  3,600 
atmospheres  for  this  purpose.  These  negative  results  find  their  explana- 
tion in  a general  property  of  gases,  first  recognized  by  Andrews  (1869), 
and  called  by  him  the  critical conditio7i  of  matter.  There  is  a temperature 
common  to  all  gases,  above  which  they  cannot  be  condensed — this  is  the 
critical  temperature.  It  was  first  observed  with  carbon  dioxide  (see  this). 
It  can  be  liquefied  at  0°  under  a pressure  of  35.4  atmospheres,  at  13° 
under  48.9  atmospheres,  and  at  30°  under  73  atmospheres.  Again,  Cag- 
niard  de  la  Tour  (1821)  showed  that  all  liquids  when  heated  above  a cer- 
tain temperature  (the  same  critical  temperature),  would  be  transformed 
into  gases  (absolute  boiling  point  of  Mendelejeff),  although  they  were 
subjected  to  intense  pressure  (in  sealed  tubes).  The  pressure  exerted  by 
the  gas  at  the  critical  temperature  (at  which  it  would  immediately  con- 
dense upon  lowering  the  temperature)  is  called  the  critical  pressure  ; the 
volume  occupied  by  the  substance  at  this  time  is  the  critical  volwiie. 

An  additional  interesting  fact  is  the  behavior  of  a liquid  when  it  is 
gradually  heated  under  a pressure  which  exceeds  its  critical  pressure. 
When  this  is  done  the  temperature  can  be  increased  to  any  extent  with- 
out producing  a noticeable  separation  into  liquid  and  gas.  At  the  critical 
temperature  the  liquid  changes  to  gas  without  becoming  heterogeneous 
and  without  sudden  increase  in  volume.  At  the  critical  temperature  and 
pressure  the  volume  of  the  gas  (or  saturated  vapor)  is  equal  to  the  volume 
of  an  equal  quantity,  by  weight,  of  the  liquid.  A difference  between  the 
gaseous  and  liquid  condition  no  longer  exists. 

The  critical  condition  also  furnishes  us  with  a criterion  for  the  distinc- 
tion between  a gas,  in  the  strictest  sense  of  the  word,  and  a vapor.  A 
vapor  by  increase  of  pressure  with  constant  temperature  may  become  a 
liquid  (or  solid),  while  this  is  not  possible  with  a gas.  A gaseous  body 
is,  according  to  this  idea,  a permanent  gas  above  the  critical  temperature, 
while  below  the  same  it  is  a condensable  vapor. 


48 


INORGANIC  CIIKMIS'J’RY. 


'To  licjiiefy  gases,  we  need  nut  only  pressure,  but  also  a definite  tem- 
perature, and  this  must  be  lower  than  the  critical.  I>y  this  means 
Cailletei  and  Pictet  ( 1 877)  succeeded  in  condensing  nearly  all  the  ptr- 
manent  gases.  Pictet  })ursued  the  method  of  Paraday,  who  had  con- 
densed various  gases  in  sealed  tubes  (see  Condensation  of  Cddorine,  p.  51). 
'J'he  gases  were  generated  in  a powerful  iron  retort  (oxygen  from  potas- 
sium chlorate  ; hydrogen  from  sodium  formate)  by  the  ai)plication  of  heat. 
They  were  then  com])ressed  under  their  own  jjressure  in  a coi)per  tube 
attached  to  the  retort.  Solid  carbon  dioxide  surrounded  the  tube,  and 
by  its  evaporation  under  the  air  i)ump  its  temperature  was  reduced  to 
— 140°.  On  opening  the  stoi)-cock  of  the  coi)per  tube,  the  liquefied 
gas  escaped  in  a stream  which  rajiidly  evaporated.  Cailletet  employed  a 
capillary  glass  tube,  provided  with  a reservoir  and  a pressure  ])ump. 
d’he  strongly  conqiressed  gas  was  cooled  by  opening  a stoj^-cock  and  [ler- 
mitting  it  to  expand  suddenly.  In  its  exj^ansion  and  in  overcoming  the 
external  pressure  it  i)erforms  work  and  there  follows  an  absorj)tion  of  an 
ai)})reciable  quantity  of  heat,  which  is  taken  from  the  gas.  This  causes  a 
partial  liquefaction  of  the  gas  in  the  form  of  a dense  cloud,  or  in  small 
drops. 

With  these  facts  as  a basis  Wroblewsky,  Olszewsky,  and  more  especially 
Dewar,  Linde,  and  Ramsay  have  succeeded  in  condensing  and  retaining 
all  gases  in  liquid  form  ; indeed  most  of  them  have  been  solidified. 
Usually  Cailletet’s  method  has  been  pursued,  care  be  taken  to  thoroughly 
cool  the  gas  before  the  sudden  decrease  in  pressure  occurs  (see  Hydrogen 
and  Air).  Pictet’s  suggestion  to  strongly  chill  a greatly  condensed  gas 
has  proved  satisfactory  for  the  condensation  of  several  permanent  gases 
to  liquid  form.  Carbon  dioxide,  liquid  ethylene,  liquid  oxygen,  and 
nitrogen,  which  vaporize  at  slight  pressure,  have  also  been  employed  to 
produce  the  reduction  in  temperature.  With  a pressure  of  10  mm.  the 
temperature  of  liquid  ethylene  is  reduced  to  — 150°,  liquid  oxygen  at 
9 mm.  to  — 211°,  and  liquid  nitrogen  to  — 225°.  Lower  temperatures 
than  these,  about  — 250°,  are  only  attainable  by  the  evaporation  of  liquid 
hydrogen. 

Liquefied  gases  can  be  preserved  quite  well  in  Dewar  bulbs  : open, 
double-walled  glass  vessels,  the  intermediate  space  of  which  is  a vacuum 
and  the  inner  wall  coated  on  its  exterior  with  a thin  layer  of  mercury. 
The  conduction  and  radiation  of  heat  are  thus  reduced  to  a minimum. 

The  critical  temperatures  (T)  and  critical  pressures  in  atmospheres  (P) 
of  the  gases  condensed  with  difficulty  are  as  follows: 


Carl)on  Dioxide,  CO^, 
I'khylene,  C.^lf,,  . . . . 
Nitric  Oxide,  NO,  . . . 
Marsh  Oas,  ClI^,  . . . 
Oxygen,  O.^,  . . . 

C’arhon  Monoxide,  CO,  . 

Nitrogen,  N.„ 

Hydrogen,  11^,  . . . . 

Air, 


T 

P 

+ 31° 

77  atmos. 

+ 10° 

51  “ 

- 93° 

71  “ 

— 82° 

55  “ 

— 118° 

50  “ 

—140° 

35  “ 

— 146° 

35  “ 

— 220° 

15  “ 

— 140° 

39  “ 

Low  temiieratiires  are  ascertained  by  means  of  a hydrogen  or  helium 


HALOGEN  GROUP— CHLORINE. 


49 


thermometer,  a thermo-electric  element  composed  of  copper  and  German 
silver,  or  a platinum  resistance  thermometer. 

The  critical  temperature,  pressure  and  volume  can  be  determined  not  only  experiment- 
ally, but  also  may  be  deduced  from  the  variations  of  the  gases  from  the  laws  of  Boyle  and 
Gay-Lussac,  by  a theory  developed  by  van  der  Waals  (1873)  (compare  Air  : measurement 
of  gases).  Students  of  chemistry  desiring  more  detailed  information  on  this  subject 
are  directed  to  the  works  of  Ostwald  : Grundriss  der  allgemeinen  Chemie,  2 Aufl. 
1890,  and  Lothar  Meyer  : Chemical  Theories,  1893,  and  also  Ostwald’ s Lehrbuch, 
vol.  I,  289  (1891). 


,,  ■ ' ' 

■ "’halogen  group. 

Chlorine,  bromine,  iodine,  and  fluorine  constitute  this  group.  These 
elements  show  a similar  chemical  deportment.  They  are  termed  halogens 
or  salt  producers,  because  by  their  direct  union  with  the  metals  salt-like 
derivatives  result.  Chlorine,  bromine,  and  iodine  will  be  discussed  first 
because  of  their  great  importance. 


1.  CHLORINE. 

Atom  ; Cl  = 35.45.  Molecule : CL  = 70.90. 

It  does  not  occur  free  in  nature.  As  it  acts  very  energetically  upon 
most  bodies  containing  hydrogen,  especially  those  of  an  organic  nature, 
to  form  new  derivatives  with  them  or  with  their  constituents,  chlorine  can 
only  be  obtained  free  in  the  transitional  state.  Its  most  important  deriv- 
ative is  sodium  chloride,  or  rock  salt,  which  is  composed  of  chlorine  and 
sodium.  The  Swedish  chemist,  Scheele,  discovered  chlorine  in  1774. 
It  was  regarded  as  a compound  body,  until  its  elementary  character  was 
established  by  Gay-Lussac  and  Thenard  in  France  (1809),  and  by  Davy 
in  England  (1810). 

Preparation. — To  obtain  free  chlorine,  heat  a mixture  of  black  oxide 
of  manganese  (MnOa)  and  hydrochloric  acid  in  a flask  (Fig.  31),  pro- 
vided with  a so-called  Welter  safety-tube  to  equalize  the  gas  pressure. 
The  escaping  gaseous  chlorine  is  washed  and  freed  from  acid  that  is 
carried  along  mechanically  by  passing  it  through  water  in  a three-necked 
Woulff  bottle,  and  then  collecting  it  over  water.  The  reaction  which 
occurs  above  is  indicated  in  the  following  equation  : 

MnO^  4-  4HCI  = MnCh  + 21120  -f  Clj. 

The  manganous  chloride  formed  dissolves  in  the  water. 

At  low  temperatures  the  mangane.se  peroxide  dissolves  in  the  hydrochloric  acid  without 
the  evolution  of  chlorine,  the  solution  being  brown  in  color  : 

Mn02  + 4IICI  = MnCq  + 2II2O. 

Chlorine  is  expelled  on  the  aj)plication  of  heat : 

MnCb  = MnCl2  -f  Cl^. 


50 


INOKflANFC  CHEMISIRV. 


The  evolution  of  tlie  chlorine  proceeds  more  regularly  if  a mixture  of  manganese  oxide, 
sodium  chloride,  and  suliduiric  acid  is  employed. 

The  hydrochloric  acid  is  formed  from  the  last  two  (together  with  sodium  acid  sul- 
phate) : 

NaCl  f II, SO,  zzzz  IICl  + NallSO,, 

and  then  acts  upon  the  manganese  peroxide.  It  is  advisable  to  use  5 parts  of  manga- 
nese peroxide  (90  j^er  cent. ),  ll  parts  of  sodium  chloride,  atid  14  parts  of  concentrated 
sulphuric  acid  diluted  with  7.5  parts  of  water  [Klason,  l>er.  23  (1890),  330]. 

'File  second  method  is  more  advantageous  for  laboratory  jnirposes  ; the  first  (action  of 
hydrochloric  acid  upon  manganese  dioxide),  however,  is  jireferred  in  jiractice,  as  it  is 
cheaper. 

d'he  resulting  manganous  chloride  (MnCl,)  is  converted  by  the  proctss  of  Weldon 
into  manganese  peroxide  (see  this). 

Other  technical  methods  for  the  preparation  of  chlorine  are  based  upon  the  liberation 
of  chlorine  from  hydrochloric  acid  or  metallic  chlorides  (particularly  calcium  chloride  and 
magnesium  chloride)  by  the  oxygen  of  the  air  at  elevated  temperatures.  Ily-products  of 


Fig.  31. 

the  soda  industry,  mother  liquors  of  sea  salt  and  tailings  from  the  .Stassfurth  mines,  all  of 
which  contain  chlorine,  are  utilized  in  this  way. 

1.  Deacon- JIurte)'  Process:  Hydrochloric  acid  mixed  with  air  is  conducted  over 
jiorous  substances  saturated  with  copper  chloride  or  sulphate  and  heated  to  370-400°. 

2.  Weldo)i- Pechiney  Process  : Magnesium  chloride  or  oxychloride  is  heated  to  1000° 
in  an  air  current  : MgCl,  | O = MgO  -|-  Cl, ; the  magnesia  is  reconverted  into  chlo- 
ride by  solution  in  hydrochloric  acid. 

3.  Solvay  Process:  Calcium  chloride  or  magnesium  chloride,  mixed  with  silica  or 
clay,  is  heated  in  an  air  current:  CaCl,  + SiO,  I O =CaOSiO.,  4-  Cl.,. 

4.  Pcychlcr-de  Wilde  Mel  hod : A mixture  of  magnesium  chloride,  manganous  chlo- 
ride and  manganc.se  sulphate  is  exposed  to  air  heated  to  about  520°  until  tlie  evolution 
of  chlorine  ceases,  when  it  is  restored  by  hydrochloric  acid  to  its  original  condition  and 
chlorine  again  liberated. 

t;.  Mond  Method : Ammonium  chloride  vapor  is  allowed  to  act  upon  magne.sium 
oxide  or  olliei'  metallic  oxides  .and  the  ])roducts  are  ammonia  and  metallic  chlorides, 
'file  latter  aie  heated  to  a high  tempc'iatuie  in  air  when  chlorine  and  oxiiles  result. 


CHLORINE. 


51 


The  electrolytic  decomposition  of  salt  solutions  is  especially  important ; 
indeed  it  is  likely  to  supplant  the  other  technical  i)rocesses  in  the  near 
future.  The  current  decomi)oses  sodium  chloride  (and  other  metallic 
chlorides)  into  its  components,  so  that  chlorine  separates  at  the  positive 
and  sodium  at  the  negative  pole.  The  sodium  at  once  acts  upon  water 
forming  sodium  hydroxide  and  hydrogen.  This  universally  important 
process  will  be  again  referred  to  in  connection  with  soda,  potassium  chlo- 
rate and  bleaching  lime.* 

An  excellent  laboratory  method  for  the  preparation  of  chlorine  con- 
sists in  allowing  dilute  hydrochloric  acid  to  act  upon  bleaching  lime  (see 
this).  The  latter  is  previously  mixed  with  burnt  gypsum  ()^  part),  and  a 
little  water  added,  when  the  mass  can  be  formed  into  cubes  or  stout  sticks 
which  are  introduced  into  a Kipp  generator  (p.  42). 

Very  pure  chlorine  may  be  obtained  by  digesting  potassium  bichro- 
mate with  concentrated  hydrochloric  acid  : 

KaCr^O^  + 14HCI  = Cr2Cl6  + 2KCI  -f  7H2O  + 3CI2. 

As  chlorine  dissolves  readily  in  cold  water  it  is  advisable  to  collect  it  over  warm  water. 
It  cannot  be  collected  over  mercury,  as  it  readily  combines  with  the  latter.  When  perfectly 
dry  chlorine  is  sought,  conduct  the  liberated  gas  through  Woulff  bottles  containing  con- 
centrated sulphuric  acid,  to  absorb  the  moisture,  then  collect  in  an  empty  upright  flask. 
As  chlorine  is  so  much  heavier  than  air  it  will  displace  the  latter. 

Physical  Properties. — Chlorine  is  a yellowish-green  gas-  (hence  its  name 
irom  yXwpoq),  with  a penetrating,  suffocating  odor.  Its  specific  gravity 
compared  with  oxygen  (02  = 32)  is  70.9  ; 
with  air  (—  i)  it  is  2.45.  A liter  of  chlo- 
rine at  0°  and  760  mm.  pressure  at  the  sea- 
level  in  the  latitude  of  45°  weighs  3.167 
grams.  At  15°  C.,  and  a pressure  of  57  at- 
mospheres (at  — 40°  C.,  under  the  ordinary 
pressure)  it  condenses  to  a yellow  liquid, 
boiling  at  — 33  6°.  By  pressure  and  cold, 
liquid  chlorine  has  been  made  applicable 
in  technical  operations  (see  Knietsch,  Ann. 

Chem.  259  (1890),  100).  The  specific 
gravity  of  liquid  chlorine  at  its  boiling 
point  is  1.5575.  At  very  low  tempera- 
tures it  solidifies  to  a yellow  crystalline 
mass,  which  remelts  at  — 102°.  The 
critical  temperature  of  chlorine  is  146°,  and  its  critical  pressure  93.5 
atmospheres.  It  may  be  condensed  in  the  following  manner  as  demon- 
strated by  Faraday:  Take  a bent  glass  tube  (Fig.  32),  introduce  into 
the  leg  closed  at  one  end  crystals  of  chlorine  hydrate  (-CI2  + SHjO,  see 
p.  52),  then  seal  the  o])en  end.  The  limb  containing  the  comi)ound  is 


Fig.  32. 


■^‘Compare  N.  Caro,  Dar.stellung  von  Cblor  unci  Salzsaure,  Tlerlin,  1893;  also  L. 
Mond,  flesdiichtliclier  Ue])erblick  der  verscbiedenen  Arten  der  Cldordarstelliing,  Cliem- 
iker-Zeitung,  1896,  928,  and  R.  Masenclever,  Die  Entwickelung  der  Sodafabrikation, 
Ber.  29  (1896),  III,  2861. 


52 


INORGANIC  CHEMISTRY. 


placed  in  a water-bath;  the  other  is  cooled  in  snow.  Uj)on  heating  the 
water  a little  above  30°  the  chlorine  hydrate  is  decomposed  into  water 
and  chlorine  gas,  which  condenses  to  a liquid  in  the  cooled  limb.  On 
reversing  the  position  of  the  limbs  and  cooling  the  one  previously 
warmed,  the  chlorine  distils  back  and  is  reabsorbed  by  the  water.  Ohar- 
coal  saturated  with  chlorine  may  be  substituted  for  the  chlorine  hydrate. 
One  volume  of  charcoal  takes  up  200  volumes  of  chlorine,  which  are 
disengaged  again  on  heating. 

One  volume  of  water,  at  20°  C.,  absorbs  2 volumes  of  chlorine;  at 
10°  C.,  2.5  volumes.  The  aqueous  solution  is  known  as  chlorine  water 
{aqua  chlorata),  and  possesses  almost  all  the  ])roperties  of  the  free  gas;  it 
is  therefore  frequently  employed  for  laboratory  uses  as  a substitute  for 
chlorine. 

Recent  investigations  show  that  chlorine  water,  by  virtue  of  the  action  of  the  chlorine 
upon  water,  contains  hydrochloric  and  hypochlorous  acids,  which  in  turn  react  upon  one 
another  with  the  production  of  chlorine  and  water,  so  that  in  accordance  with  the  strength 
of  the  chlorine  water  and  external  prevailing  conditions  equilibrium  occurs  among  these 
four  substances  (p.  29).  This  may  be  expressed  as  follows  : 

Cb  + H-P  HCl  + HCIO. 

The  yellow,  scale-like  crystals  of  chlorine  hydrate  (Cb  -j-  8H2O)  sepa- 
rate when  water  saturated  with  the  gas  is  cooled  below  0°.  This  com- 
pound is  regarded  as  one  of  chlorine  with  water.  At  ordinary  tempera- 
tures it  decomposes  into  water  and  chlorine. 

Chemical  Properties. — When  thin  sheet  copper  (false  gold  leaf),  or, 
better,  pulverized  antimony  or  arsenic,  is  thrown  into  a vessel  filled  with 
dry  chlorine,  it  will  burn  with  a bright  light;  a piece  of  phosphorus  will 
also  inflame  in  an  atmosphere  of  the  gas.  Perfectly  dry  chlorine  does 
not  manifest  this  energetic  combination-tendency. 

Chlorine  unites  just  as  energetically  with  hydrogen.  A mixture  of 
equal  volumes  of  the  gases  obtained  by  the  electrolysis  of  hydrochloric  acid 
combines  in  direct  sunlight  with  violent  explosion.  In  diffused  sun- 
light the  action  is  only  gradual ; in  the  dark  it  does  not  occur  if  perfectly 
dry  gases  are  employed  (Baker).  The  course  of  this  reaction  and  the 
laws  of  the  chemical  action  of  light  have  been  ably  discussed  in  a series 
of  celebrated  papers  by  Bunsen  and  Roscoe  [Pogg.  Ann.  d.  Physik,  Bd. 
100,  lOi,  108  (1857-1859)].  Chlorine  also  manifests  great  affinity  for 
hydrogen  derivatives,  most  of  them  being  so  decomposed  by  the  chlorine 
that  hydrogen  is  removed  from  them,  and  hydrochloric  acid  is  formed. 
Thus  water  is  decomposed,  especially  in  sunlight,  by  chlorine  into  hydro- 
chloric acid  and  oxygen  : 

11,0  + CI2  2HCI  + O. 

If  a glass  cylinder  be  filled  with  and  inverted  over  chlorine  water  and 
exj)osed  to  direct  sunlight,  a gas  will  be  evolved,  and  will  collect  in  the 
ui)i)er  |)ortion  of  the  vessel  ; this  isoxygen.  In  diffused  light  the  decom- 
position will  not  be  so  rapid  ; it  is  hastened  by  heat.  Chlorine  oxygen 
acids,  IK'IO,  IICK).,  fsec  above),  are  also  jiroduced. 

Chlorine  alters  the  hydrocarbons,  in  that  it  abstracts  hydrogen.  The 


BROMINE. 


53 


reaction  is  sometimes  so  violent  that  carbon  is  separated  in  a free  condi- 
tion. A piece  of  tissue  paper  saturated  with  freshly  distilled  turpentine 
oil,  and  introduced  into  a dry  chlorine  atmosphere,  is  immediately  car- 
bonized. An  ignited  wax  taper  immersed  in  chlorine  burns  with  a smoky 
flame,  with  separation  of  carbon. 

The  organic  dyestuffs  (containing  carbon  and  hydrogen)  are  decolor- 
ized by  moist  chlorine  gas.  The  same  occurs  with  the  dark-blue  solutions 
of  indigo  and  litmus ; colored  flowers  are  rapidly  bleached  by  it.  On  this 
princii)le  depends  the  application  of  chlorine  in  bleaching,  and  the  de- 
struction of  decaying  matter  and  miasmata  in  chlorine  disinfection  are 
dependent  upon  this  reaction  (see  Bleaching  Lime). 

The  bleaching  action  of  chlorine  is  mostly  influenced  by  the  presence  of  water.  It 
probably  depends  on  the  oxidizing  action  of  the  oxygen  liberated  by  the  chlorine  or  upon 
the  chlorine  oxy-acids  found  when  chlorine  and  water  react  (see  p.  52).  This  property, 
free  oxygen  does  not  possess  ; it  does,  however,  very  probably  belong  to  that  which  is  in 
the  act  of  fonning, — of  becoming  free.  We  will  learn,  later,  that  many  other  elements, 
at  the  moment  of  their  birth  {in  statu  nascendi),  act  more  energetically  than  when  in 
the  free  condition. 


2.  BROMINE. 

Atom  : Br  = 79.96.  Molecule : Br2  = 159.92. 

Bromine,  the  perfect  analogue  of  chlorine,  was  discovered  by  Balard, 
in  1826.  It  is  not  found  free  in  nature  for  the  same  reasons  which 
were  given  under  chlorine.  It  occurs  in  sea-water  as  sodium  bromide, 
accompanied  by  sodium  chloride,  but  in  much  smaller  quantity  than  the 
latter  (especially  in  the  water  of  the  Dead  Sea),  and  in  many  salt  springs, 
as  at  Kreuznach  and  in  Hall.  When  sea-water  or  other  salt  water  is 
evaporated,  sodium  chloride  first  separates;  in  the  mother  liquor,  among 
other  soluble  salts,  are  found  sodium  and  magnesium  bromides.  Bromine 
is  found  in  greatest  abundance  in  the  upper  layers  of  the  rock-salt  deposits 
of  Stassfurth,  near  Magdeburg,  where  it  exists  in  the  form  of  bromides 
together  with  other  salts.  At  present,  large  quantities  of  bromine  are 
obtained  in  America.  The  method  of  its  preparation  is  similar  to  that 
employed  under  chlorine.  A mixture  of  manganese  dioxide  and  sodium 
bromide  is  warmed  with  sulphuric  acid  : 

MnOg  + 2NaBr  -f  2H2SO^  = MnS04  -|-  Na^SO^  -|-  Br2  -f  2H2O. 

The  bromine  condenses  in  the  well-cooled  receivers.  When  free  chlo- 
rine is  conducted  into  an  aqueous  solution  of  sodium  bromide,  bromine 
separates.  This  is  the  method  adopted  at  Stassfurth  and  in  America  in 
preparing  bromine  from  the  magnesium  bromide  saline  liquors : 

MgBr2  -f  CI2  = MgCb  -f  Br2. 

Bromine  is  a heavy,  reddish-brown  liquid,  with  an  exceedingly  pene- 
trating, chlorine-like  odor  (hence  the  name  Bromine,  from  [ipcbij.oc;,  stench). 
When  bromine  is  strongly  cooled  it  solidifies  to  a dark-brown  mass,  of 
delicate  needles  with  a slightly  black  metallic  luster,  which  remelt  at 
— 7.3°.  Liquid  bromine  at  0°  has  the  sjiecific  gravity  3. 18  (water  = i)  ; 
it  is  very  volatile,  forming  dark-brown  vapors  at  the  ordinary  tempera- 


54 


INORGANIC  CHEMISTRY. 


lure,  and  boils  at  about  6o°,  clianging  at  the  same  time  into  a dark 
brown-red  va^xir.  Its  density  eijiials  159.9  (oxygen  32),  (jr  5.52  (air 
= 1). 

One  part  of  bromine  dissolves  in  about  35  ])arts  of  water  of  medium 
temjierature;  it  is  therefore  more  soluble  in  water  than  chlorine.  Cooled 
below  4°  C.,  the  hydrate  ( I Ir.^ -f  roll./))  crystallizes  out:  this  is  analogous 
to  the  chlorine  hydrate.  It  is  decomi)osed  at  6.2°.  J>romine  dissolves 
with  ease  in  alcohol,  and  esi)ecially  in  ether,  chloroform  and  bisulphide. 

Chemically,  bromine  is  extremely  like  chlorine,  combining  directly 
with  most  metals  to  form  bromides;  but  it  possesses  a weaker  affinity 
than  chlorine,  and  is  liberated  by  the  latter  from  its  compounds.  With 
hydrogen  it  only  combines  on  warming,  not  in  sunlight.  Ujion  hydro- 
carbons it  acts  like  chlorine,  withdrawing  hydrogen  from  them,  llromine 
water  gives  starch-paste  an  orange  color. 

Solid  bromine  is  a mixture  of  licpiid  bromine  with  silicious  earth.  It 
occurs  in  trade  pressed  into  cubes  or  sticks  and  is  used  for  disinfecting 
purposes.  Bromine  and  its  salts  are  api)lied  in  medicine,  in  photography 
and  in  the  manufacture  of  aniline  colors. 


3.  IODINE.  . \ 

Atom  : I = 126.85.  Molecule  : I2  = 253.70. 

Iodine,  as  well  as  bromine,  occurs  in  combination  with  sodium,  in  sea- 
water and  some  mineral  springs,  especially  at  Hall,  in  Austria,  and  the 
Adelheit  spring  in  Bavaria.  In  these  springs  the  iodine  can  easily  be  de- 
tected; in  sea-water  it  is,  however,  only  present  in  such  minute  quantity 
that  its  separation,  practically,  is  disadvantageous.  Sea  algae  absorb  iodine 
compounds  from  the  water,  and  these  are  then  thrown  by  the  tide  on  various 
coasts,  where  they  are  burned,  yielding  an  ash  (known  as  kelp  in  Scotland, 
as  varec  in  Normandy)  which  is  the  jirincipal  source  for  the  manufacture 
of  iodine.  It  was  in  this  ash  that  the  element  was  accidentally  discovered 
l)y  Courtois  in  1811;  in  1813,  it  was  investigated  by  Davy  and  Gay- 
Lussac,  and  its  elementary  character  established.  To  obtain  the  iodine, 
the  ash  is  treated  with  water,  the  solution  concentrated,  and  separated 
from  the  chlorides  of  sodium  and  potassium  which  have  crystallized  out, 
the  final  liquors,  in  which  the  readily  soluble  iodides  have  accumulated, 
being  then  distilled  with  manganese  dioxide  and  sulphuric  acid.  Iodine, 
therefore,  is  set  free  from  its  compounds  in  the  same  manner  as  chlorine 
and  bromine.  It  is  more  convenient,  however,  to  pass  chlorine  (or 
better,  nitrous  acid)  through  a solution  of  the  iodides,  when  all  the  iodine 
will  separate : 

KI  4-  Cl  KCl  4- 1, 

or  the  iodide  solution  is  digested  with  ferric  chloride,  when  the  latter  is 
reduced  to  ferrous  chloride  : 

Nal  1-  FeCly  =1-- 1 + NaCl  |-  FeCl.,. 

The  grayish-black  jiowder  thus  liberated  is  collected  on  a filter,  dried, 
and  then  sublimed. 


IODINE. 


55 


In  recent  years  the  greatest  quantities  of  iodine  have  ])een  obtained 
from  the  mother  liquors  of  crude  Chile  saltpeter  (NaNO;,).  The  iodine  is 
present  in  this  salt  as  sodium  iodate  (NalOa),  from  which  it  is  set  free  by 
nitrous  or  sulphurous  acid  : 

aNalOg  + SH.SOg  = 2I  -f  2NaHS04  + SH.SO^  + Hp. 

An  interesting  occurrence  of  iodine  is  that  observed  by  Baumann  in 
1895  thyroid  (thyroidea,  from  Oopeoq,  shield,  and  e'ldof;,  form)  gland, 

the  active  constituent  of  which,  containing  9 per  cent,  of  iodine,  is  mixed 
with  milk-sugar  and  sold  in  the  markets  under  the  name  thyroidin. 

Iodine  is  a gray-black  solid,  subliming  in  large  rhombic  crystals,  pos- 
sessing strong  metallic  luster.  It  has  a peculiar  odor,  reminding  one 
somewhat  of  that  of  chlorine;  it  stains  the  skin  brown,  and  is  corrosive, 
although  not  as  strongly  so  as  bromine.  Its  specific  gravity  is  4.95  at 
15°.  It  fuses  at  114°  to  a dark-brown  liquid,  and  boils  near  184.3°, 
passing  at  the  same  time  into  a dark-violet  vapor  (hence  the  name  Iodine, 
suggested  by  Gay-Lussac,  from  loetdrjq,  violet-like). 

The  vapor  density  of  iodine  equals  8.7  up  to  600°  C.  (air=  i)  or  254  (O2—  32),  cor- 
responding to  the  molecular  weight  I2  = 253.7.  Above  600°  the  vapor  density  gradually 
diminishes,  and  at  about  1500°  it  is  only  half  the  original.  This  is  explained  by  the 
gradual  decomposition  (see  Dissociation  of  Water)  of  the  normal  diatomic  molecule  L 
into  the  free  atoms  I T I.  In  like  manner  the  bromine  molecules  Br2  suffer  a separation 
into  the  free  atoms.  The  dissociation  of  bromine  vapor  (diluted  with  ii  volumes  of  nitro- 
gen) commences  at  about  1000°  and  is  complete  at  1600°.  The  vapor  density  of  chlorine 
is  still  normal  at  1200°,  and  it  is  only  at  1400°  that  it  sustains  a slight  diminution.  Oxy- 
gen and  nitrogen  on  the  contrary  show  no  alteration  in  their  vapor  density  even  at  1690° 
(C.  Danger  and  V.  Meyer). 


Iodine  is  very  slightly  soluble  in  water  (i  : 3600),  more  readily  in 
alcohol  (^Tinctura  iodi'),  very  easily  in  aqueous  potassium  iodide,  in  ether, 
in  chloroform  and  in  carbon  bisulphide  (10  c.c.  of  the  latter  will  dissolve 
1.85  grams  of  iodine  at  18°),  the  last  two  assuming  a deep  red-violet  color 
in  consequence.  It  affords  a particularly  beautiful  crystallization,  consist- 
ing of  forms  of  the  rhombic  system,  when  it  separates  from  a solution  of 
glacial  acetic  acid. 

In  chemical  deportment  iodine  closely  resembles  bromine  and  chlorine ; 
it  possesses,  however,  weaker  affinities  and  for  this  reason  is  liberated  from 
its  compounds  by  those  elements.  With  the  metals  it  usually  combines 
only  when  warmed  ; with  hydrogen  it  does  not  combine  directly,  and  it 
does  not  remove  it  from  its  carbon  compounds. 

The  deep-blue  color  it  imparts  to  starch  is  characteristic  of  iodine. 
On  adding  starch-])aste  to  the  solution  of  an  iodide,  and  following  this 
with  a few  drops  of  chlorine  water,  the  paste  will  immediately  be  colored 
a dark  blue  by  the  separated  iodine.  This  reaction  serves  to  detect  the 
smallest  quantity  of  it. 

Iodine  is  largely  employed  in  medicine,  photography,  and  in  the 
preparation  of  aniline  colors. 


56 


INORGANIC  CHKMIS'IKY. 


4.  FLUORINE. 

Atom  ; r'l  = 19.  Molecule  : l''l2  = 38. 

Fluorine  has  such  great  affinity  for  nearly  all  substances,  that  des])ite 
numerous  efforts  it  has  only  recently  been  possible  to  obtain  it  in  a free 
condition.  It  was  Moissan  (1886)  who  electrolyzed  anhydrous,  strongly 
chilled  hydrofluoric  acid  to  which  some  sodium  lluoride  had  been  added, 
and  obtained  fluorine  as  a slightly  greenish-yellow  colored  gas.  The  de- 
composing vessel  was  a j)latinum  tube  with  sto])pers  of  fluorspar.  'J'he 
])ositive  pole,  at  which  the  fluorine  ai)peared,  consisted  of  an  alloy  of 
platinum  and  iridium.  Dewar  and  Moissan  have  very  recently  succeeded 
in  liquefying  fluorine  at  — 187°  by  the  use  of  boiling  oxygen  as  the  refrig- 
erating substance.  Under  these  conditions  it  loses  its  chemical  activity 
almost  completely,  does  not  attack  glass,  iodine,  sulphur,  or  metals. 
It  is  only  the  affinity  for  hydrogen  which  remains;  at  least  benzene  is 
entirely  decomposed  with  production  of  flame  by  fluorine. 

Fluorspar  (CaFla)  is  the  most  important  fluorine  derivative.  It  is  most 
frequently  employed  in  the  preparation  of  other  fluorine  compounds. 

Fluorine  combines  with  hydrogen,  iodine,  sulphur,  silicon,  boron,  pow- 
dered arsenic  and  antimony,  finely  divided  iron  and  manganese  even  in 
the  dark ; organic  substances,  like  oil  of  turpentine,  alcohol  and  cork 
burn  in  the  gas.  The  metals,  with  the  exception  of  gold  and  platinum, 
are  energetically  attacked  by  fluorine.  The  latter  decomposes  water  with 
the  production  of  hydrofluoric  acid  and  ozonized  oxygen,  and  liberates 
chlorine,  bromine  and  iodine  from  their  metallic  derivatives.  The 
density  of  fluorine  referred  to  hydrogen  equals  1.32  (1.26  found). 

These  four  elements,  fluorine,  chlorine,  bromine  and  iodine,  exhibit 
gradual  differences  in  their  properties;  and,  what  is  remarkable,  this  gra- 
dation stands  in  direct  relation  to  the  specific  gravity  of  the  elements  in 
the  state  of  gas  or  vapor. 

FI  Cl  Br  I 

Specific  gravity,  19  35-45  79- 9^  126.85 

A simultaneous  condensation  of  matter  occurs  with  the  increase  of  spe- 
cific gravity.  This  expresses  itself  in  the  diminished  volatility.  Fluorine 
is  a gas  down  to  — 187°  ; chlorine  can  readily  be  condensed  to  a liquid ; 
bromide  is  a liquid  at  ordinary  temperatures,  and  iodine  is  a solid. 
Other  ])hysical  properties,  as  seen  in  the  following  table,  are  also  in 
accord  with  the  preceding: 


Fluorine. 

Chlorine. 

Bromine. 

Iodine. 

Fusing  point, 

— 102° 

— 72° 

4 114° 

Foiling  ]X)inl, 

-187°  (?) 

- 33° 

+ 60° 

4184° 

Specific  gravity  in  licpiid  .... 

1. 14 

1.47 

3.18 

or  .solid  condition,  .... 

4-95 

Color, 

Green  - yel- 
low 

Yellow- 

green 

Frown 

Flack- 

violet 

HYDROGEN  CHLORIDE. 


57 


Just  such  a gradation,  as  we  have  seen,  is  observed  in  the  chenaical 
affinities  of  these  four  elements  for  the  metals  and  hydrogen  ; fluorine  is 
the  most  energetic,  iodine  the  least.  Iodine  is  therefore  separated  from 
its  soluble  metallic  and  hydrogen  compounds  by  the  other  three,  bromine 
by  chlorine  and  fluorine,  and  chlorine  by  fluorine  (p.  56).  We  shall  dis- 
cover, later,  that  the  halogens  are  displaced  in  exactly  the  reverse  order 
of  their  oxygen  compounds ; that  iodine  has  the  greatest  and  chlorine  the 
feeblest  affinity  for  oxygen.  Oxygen  derivatives  of  fluorine  are  not 
known. 


COMPOUNDS  OF  THE  HALOGENS  WITH 
HYDROGEN. 

The  halogens  form  gaseous  acids,  readily  soluble  in  water,  with  hydro- 
gen : the  halogen  hydrides  or  haloid  acids.  But  one  such  compound  is 
known  for  each  halogen. 


1.  HYDROGEN  CHLORIDE. 

HCl  = 3646. 

The  direct  union  of  chlorine  with  hydrogen  takes  place  through  the 
agency  of  heat,  and  by  the  action  of  direct  sunlight  or  other  chemically 
active  rays  (magnesium  light) ; in  diffused  light  the  action  is  only 
gradual,  and  does  not  occur  at  all  in  the  dark.  When  both  gases  are 
perfectly  dry  they  do  not  react  in  direct  sunlight.  On  introducing  a 
flame  of  hydrogen  ignited  in  the  air  into  a cylinder  filled  with  chlorine 
it  will  continue  to  burn  in  the  latter  with  the  production  of  hydrogen 
chloride.  The  opposite,  the  combustion  of  chlorine  in  an  atmosphere  of 
hydrogen,  may  be  shown  easily  by  the  following  experiment  : An 
inverted  cylinder  is  filled  with  hydrogen  by  displacement,  the  gas  is 
ignited  at  the  mouth,  and  a tube  immediately  introduced  which  will  con- 
duct dry  chlorine  into  the  cylinder.  The  burning  hydrogen  will  inflame 
the  chlorine,  which  will  continue  to  burn  in  the  former.  From  these 
experiments,  we  perceive  that  combustion  is  a phenomenon  which  ac- 
companies a chemical  change ; in  this  instance  the  union  of  hydrogen 
with  chlorine;  if  hydrogen  is  combustible  in  chlorine  (or  air),  so, 
inversely,  is  chlorine  (or  air)  combustible  in  hydrogen  for  the  same 
reason.  By  the  term  combustion,  in  chemistry,  is  understood  every 
chemical  union  of  a body  with  a gas,  which  is  accompanied  by  the  phe- 
nomenon of  light. 

A mixture  of  ecjual  volumes  of  chlorine  and  hydrogen  is  called  chlor- 
detonating gas ; it  explodes  with  very  great  violence  under  the  conditions 
given  above  for  the  union  of  the  gases.  The  product  is  gaseous  hydrogen 
chloride. 

The  formation  of  the  latter  compound  succeeds  best  by  allowing  sul- 


58 


INORGANIC  CHEMISTRY. 


l)hiiric  acid  to  act  upon  sodium  chloride  when  solid  sodium  bisulphate 
and  hydrogen  chloride  gas  will  result  : 

NaCl  + 1 1, SO,  -r  NallSO,  + IICl. 

Pour  over  5 parts  sodium  cliloride,  9 })arts  of  concentrated  sulphuric  acid,  diluted  with 
water  (2  parts),  and  gently  warm  the  tlask  on  a sand-bath  (Fig.  33).  'I'he  escaping 
hydrogen  chloride  is  conducted  through  a Woulff  bottle  containing  .sulphuric  acid  or 
through  the  cylinder  ll  ((died  with  pumice-stone  saturated  with  sul[)huric  acid),  intended 
to  free  it  from  all  moisture,  and  atlerward  collected  over  mercury  in  the  cylinder  C. 

At  a red  heat,  the  acid  sodium  sulphate  reacts  anew  upon  the  sodium  chloride  with  the 
formation  of  neutral  sodium  sulphate  and  hydrochloric  acid: 

NallSO,  -f  NaCl  = IICl  -f  Na,SO,. 

These  are  the  reactions  by  which  hydrochloric  acid,  in  the  preparation  of  sodium  sulphate 
by  the  Le  Blanc  soda  process,  is  technically  made.  As  this  method,  however,  is  being 


Fig.  33. 

replaced  by  the  electrolytic  method,  in  which  not  hydrochloric  acid,  but  chlorine  is  evolved, 
it  seems  very  important  that  the  latter  should  be  converted  by  some  convenient  procedure 
into  hydrochloric  acid,  always  of  the  greatest  importance  technically.  Lorenz  contends 
that  this  may  be  accomplished  by  conducting  the  chlorine  together  with  steam  over  ignited 
coke,  when  the  chief  products  will  be  hydrochloric  acid  and  carbon  dioxide  (CO.^)  : 

2CI,  -f  2ll,0  -(-  C ==  4HCI  -f  CO.,. 

[Comj)are  Per.  30  (1897),  l,  347]. 

'I’lu!  suggestion  of  Davy  to  allow  .sulphuric  acid  to  act  upon  pieces  of  ammonium 
chloride  will  give  a regular  current  of  hydrogen  chloride: 

Nil, Cl  + II.,SO,,  = (NII,)IISO,  f IICl. 

A Norblad  generator  as  modilied  by  Kreussler  (Fig.  34)  answers  well  for  this  purpose. 


HYDROGEN  CHLORIDE. 


59 


Physical  Properties. — Hydrogen  chloride  is  a colorless  gas,  with  a 
suffocating  odor.  In  moist  air  it  forms  dense  clouds  as  it  combines  with 
the  aqueous  vai)or  to  form  hydrochloric  acid.  Its  critical  temperature  is 
about  -1-52.3°,  and  the  critical  pressure  86  atmospheres,  /.  e.,  for  its  con- 
densation at  the  tenqjerature  just  given  it  requires  a pressure  of  86  atmos- 
pheres. There  is  no  i^ressure  which  will  condense  it 
above  this  temperature.  Liquid  hydrogen  chloride  is 
colorless,  has  a specific  gravity  at  15°  of  o 83,  and  freezes 
at  — 115-7°  to  a white  crystalline  mass,  which  begins 
to  melt  at  — 112.5°.  boils  at  — 80.3°  under  the 
ordinary  pressure. 

The  specific  gravity  (density)  of  the  gas  is  36.46  (O2  = 

32),  or  1.26  (air=  i).  One  liter  of  it  weighs  1.6285  grams 
at  0°,  760  mm.  pressure,  in  latitude  45°  and  at  sea-level. 

Hydrogen  chloride  possesses  an  acid  taste,  and  colors 
in  the  presence  of  water  blue  litmus-paper  red ; it  is, 
therefore,  an  acid,  and  has  received  the  name  hydro- 
chloric acid  gas.  It  dissolves  very  readily  in  water,  and 
on  that  account  cannot  be  collected  over  it.  One  volume 
of  water  at  0°  C.  and  760  mm.  pressure  dissolves  505 
volumes,  and  at  ordinary  temperatures  about  450  volumes 
of  the  gas.  At  higher  pressure  water  dissolves  more 
hydrochloric  acid,  and  at  low  pressures  less  (compare 
Carbon  Dioxide,  and  also  Solutions).  This  great  solu- 
bility is  very  nicely  illustrated  by  filling  a long  glass 
cylinder  with  the  gas  and  then  just  dipping  its  open  end 
into  water  ; the  latter  rushes  up  into  the  vessel  rapidly  (as 
into  a vacuum),  as  it  quickly  absorbs  the  gas.  The  aque- 
ous solution  of  hydrogen  chloride  is  commonly  known 
as  muriatic  or  hydrochloric  acid  (^Acidiim  hydrochIo7'icu7)i).  For  its  prep- 
aration the  gas  is  passed  through  a series  of  Woulff  bottles  (Fig.  35)  con- 
taining water.  The  small  bottle  B,  in  which  there  is  but  little  water, 
serves  to  wash  the  gas — free  it  of  any  mechanically  admixed  sulphuric 
acid.  The  same  apparatus  may  be  employed  in  the  manufacture  of  chlor- 
ine water,  and  is  generally  used  in  the  saturation  of  liquids  with  gases. 

A solution  saturated  at  15°  C.,  contains  about  42.9  per  cent,  hydrogen 
chloride,  has  a specific  gravity  of  1.212,  and  fumes  in  the  air.  On  the 
a])plication  of  heat,  the  gas  again  escapes,  and  the  temperature  of  the 
liquid  rises  to  110°  C.,  when  a liquid  distils  over,  containing  20.24  ])ei- 
cent,  of  hydrogen  chloride,  having  a specific  gravity  of  1. 104  and  corre 
sponds  a])proximately  to  the  formula  HCl  -f-  8H.2O.  The  composition  of 
the  distillate  varies  somewhat  with  the  pressure.  A dilute  acid,  upon 
distillation,  loses  water,  until  finally  that  boiling  at  110°  C.  ])asses  over. 
On  conducting  hydrogen  chloride  into  concentrated  hydrochloric  acid 
cooled  to — 22°,  crystals  of  the  formula  HCl  -f  2H.2O  separate;  these 
fuse  at  — 18°  and  then  decompose. 

Hydrochloric  acid  finds  an  extensive  industrial  application,  and  is 
obtained  in  large  quantities,  as  a by-product,  in  the  soda  manufacture 
(Le  Blanc  process). 


Fig.  34. 


6o 


INORGANIC  CHEMISTRY. 


CJiemical  Properiies  of  Jlydroi^cn  Chloride. — Acids — Bases — Salts. — 
Hydrogen  chloride  is  a very  stable  compound  ; it  is  only  l)eyond  1500'’ 
that  it  sustains  a partial  decomposition  (see  Dissociation  of  Water).  Its 
com])Osition  is  readily  established  fpiantitatively  in  the  following  way : 
Pass  hydrochloric  acid  gas  over  a piece  of  sodium  or  iiotassium  heated  in 
a glass  tube,  and  hydrogen  will  escaiie  from  the  latter:  2Na  -f-  21101  — 
2NaCl  -f-  Il'i-  If,  on  the  other  hand,  manganese  ])eroxide  be  heated  in 
it,  chlorine  will  be  disengaged  : MnO.^ -[- 41101  = IVInOlj -f  2H2()-}-Ol2. 
If  the  electric  current  be  allowed  to  act  upon  an  a(pieous  solution  of 
hydrochloric  acid  the  latter  will  be  decomposed  so  that  chlorine  se])arates 
at  the  electro-positive  and  hydrogen  at  the  electro-negative  j)ole  (j).  77). 
Hydrogen  chloride,  as  well  as  its  solution,  possesses  all  the  properties  of 


acids,  and  can  well  figure  as  a prototype  of  these;  it  tastes  intensely  sour, 
reddens  blue  litmus-pa})er,  and  saturates  the  bases  (oxides  and  hydrox- 
ides, i.  e.,  which  bodies  impart  a blue  color  to  red  litmus-paper)  form- 
ing chlorides.  There  are  some  bases,  the  alkalies  and  alkaline  earths, 
which  are  soluble  in  water.  These  solutions  react  basic,  alkaline,  i.  e., 
red  litmus  is  colored  blue  by  them.  If  we  add  hydrochloric  acid  to  a 
solution  of  a base,  e.  sodium  hydroxide,  until  the  reaction  is  neutral, 
we  will  obtain  (besides  water)  a neutral  compound — sodium  chloride, 
which  remains  in  a crystalline  form  when  the  solution  is  evaporated: 

NaOir  I IICI  NaCI  1 II2O. 

Sodium  Sodium 

hydroxide.  chloride. 


HYDROGEN  BROMIDE. 


6l 

Hydrogen  bromide,  iodide  and  fluoride  deport  themselves  similarly  to 
hydrogen  chloride.  These  halogen  compounds  of  hydrogen  are  termed 
haloid  acids,  to  distinguish  them  from  those  which,  in  addition  to  hydrogen, 
contain  oxygen,  hence  called  oxygen  acids.  The  latter  conduct  themselves 
like  the  former,  and  saturate  bases,  forming  salts  and  water  : 


KOH  + IINO3  ^ + H^O. 

Potassium  Nitric  Potassium  Water, 

hydroxide,  acid.  nitrate. 


In  the  same  manner  the  acids  act  upon  the  basic  oxides,  to  form  salt; 
and  water : 


ZnO  + 2HCI  ==  ZnCL  4-  H^O. 
Zinc  Zinc 

oxide.  chloride. 


ZnO  + 2HNO3  = ZnfNOg)^  + H^O. 
Zinc  Zinc 

oxide.  nitrate. 


Usually  when  acids  act  upon  metals,  the  hydrogen  of  the  former  is 
directly  displaced  ; salts  and  free  hydrogen  are  produced.  Thus,  by  the 
action  of  hydrochloric  acid  upon  sodium,  its  chloride  and  hydrogen 
result : 

HCl  + Na  = NaCl  + H ; 


and  when  zinc  and  hydrochloric  acid  react,  zinc  chloride  and  water 
(see  p.  41) : 

2HCI  + Zn  = ZnCb  + H2. 


From  the  examples  cited  it  is  manifest  that  acids  are  hydrogen  com- 
pounds which  yield  salts,  by  the  replacement  of  their  hydrogen  by  metals 
(by  the  action  of  metallic  oxides,  hydroxides,  and  by  the  free  metals). 
The  metallic  oxides  and  hydroxides  like*  sodium  hydroxide,  capable  of 
forming  water  and  salts  by  the  saturation  of  acids,  are  called  bases. 
Finally,  by  the  term  salts,  we  understand  such  compounds  as  are  analogous 
to  sodium  chloride,  and  are  formed  by  the  mutual  action  of  bases  and 
acids  with  the  exit  of  water.  Salts  are  distinguished  as  haloid  salts  a.nd 
oxygen  salts.  The  first  have  no  oxygen,  and  arise  in  the  direct  union  of 
the  halogens  with  the  metals: 

Na+  Cl  =NaCl. 

Zn  -j-  CI2  = ZnCI^. 


2.  HYDROGEN  BROMIDE. 

HBr  = 80.97. 

Hydrogen  bromide  is  perfectly  similar  to  the  corresponding  chlorine 
compound.  As  there  is  but  slight  affinity  between  bromine  and  hydrogen 
their  direct  union  will  only  occur  at  a red  heat  or  in  the  presence  of 
])latinum  sponge  (see  p.  46).  Like  hydrogen  chloride,  hydrogen  bro- 
mide can  be  obtained  by  the  action  of  some  acids,  e.  g , phosphoric  acid, 
upon  bromides;  concentrated  sulphuric  acid  would  not  answer  as  the  re- 


62 


INORHANIC  CHEMIS'I’RY. 


suiting  hydrogen  bromide  is  apin  i)artly  decomposed  by  it.  Ordinarily 
it  is  prepared  by  the  action  of  phosi)horus  tribromide  upon  water  ; 

Vlh,  + 3II/)  ll3P()3  -p  3iip,r. 

PIiosi)li()nis  I’hospliorous 

tribroniide.  acid. 

Place  water  (i  part)  in  a flask  (Fig.  36),  gradually  admit  through  the 
funnel,  supplied  with  a stop-cock,  the  licjuid  j)h()S[)horus  tribromide  (3 
parts),  and  warm  gently,  'bhe  escaping  gas  is  collected  over  mercury  or 
conducted  into  water.  To  free  it  perfectly  from  accom])anying  ])hos- 
])horus  bromide  vapors  it  is  passed  through  water  (the  U-shai)ed  tube  in 
Fig.  36  contains  pieces  of  pumice-stone  or  glass  beads,  which  are  moist- 
ened with  water). 

Instead  of  employing  prepared  phosidiorus  bromide,  we  may  let  bromine  vapors  act 

u[)on  (red)  phosphorus,  d'hismay  be  done 
by  ])ouring  water  (2  parts)  over  the  phos- 
phorus ])laced  in  a flask;  bromine  ( 10  parts) 
is  added  gradually  while  cooling  and  heat  is 
then  applied.  To  free  the  hydrogen  bro- 
mide gas  from  the  bromitie  carried  along 
mechanically,  conduct  it  through  a tube 
containing  glass  wool  and  moist  red  phos- 
phorus. 

Claseous  hydrobromic  acid  can  also  be 
prepared  by  allowing  bromine  to  act  upon 
crude  anthracene.  The  resulting  hydro- 
bromic acid  gas  is  freed  from  the  accom- 
panying bromine  by  pa.ssing  it  through  a 
tube  filled  with  anthracene. 

To  obtain  an  aqueous  solution  of  the 
gas,  pour  15  parts  of  water  over  i part  of 
red  phosphorus,  and  then  add  bromine  (10 
parts)  drop  by  drop.  Finally  the  .solution 
is  heated,  filtered,  and  distilled.  Bromides 
(sodium  bromide,  potassium  bromide)  yield 
hydrogen  bromide  by  distillation  with  dilute 
Fig.  36.  sulphuric  acid  in  the  presence  of  phosphorus. 

Hydrogen  bromide  is  a colorless  gas,  fuming  strongly  in  the  air. 
Under  great  pressure  it  is  condensed  to  a liquid,  solidifying  at  — 120°, 
melting  at  — 87°,  and  boiling  at  — 73°.  Its  density  is  80.97  (O^  = 32)  or 
2.79  (air  = i). 

In  water  the  gas  is  very  readily  soluble,  its  solution  saturated  at  0° 
having  a siiecific  gravity  of  1.78,  and  containing  82  ]ier  cent,  of  hydro- 
gen bromide.  Its  composition  closely  approximates  the  formula  HBr  -|- 
1 1/)  ; at  15°  it  contains  49.8  per  cent,  of  acid  and  has  the  specific  gravity 
1.5  15.  At  125°  a solution  distils  over,  containing  48.2  per  cent,  of  hy- 
drogen bromide;  its  composition  corresponds  very  nearly  to  the  formula 
1 1 l>r  5 1 ; its  specific  gravity  is  1.49  at  14°  C. 

On  conducting  hydrogen  bromide  into  a solution  of  the  same  cooled 
to  — 20°,  crystals  of  the  formula  Hllr  -|-  2ll./)  separate  and  melt  at 
— i Chemically,  hydrogen  bromide  is  the  jierfect  analogue  of  hydro- 
gen chloride  ; it  is,  however,  less  stable,  and  suffers  a partial  decomposi- 
tion at  800°  C. 


HYDROGEN  IODIDE. 


63 


3.  HYDROGEN  IODIDE. 

HI  = 127.86. 

The  attraction  of  iodine  for  hydrogen  i.s  very  slight.  The  two  elements 
combine  at  higher  temperatures,  between  400  and  500°,  very  incom- 
pletely with  one  another,  because  at  this  point  hydrogen  iodide  is 
partly  resolved  into  its  elements  (see  below).  Their  union  is  more 
complete  if  both  elements,  in  the  form  of  vapor,  are  conducted  over 
heated  platinum  sponge.  It  cannot  be  obtained  by  acting  upon  iodides 
with  sulphuric  acid,  because  the  resulting  hydrogen  iodide  decomposes 
more  easily  than  the  bromide.  It  is  formed,  however,  similarly  to  the 
latter,  by  acting  on  phosphorus  iodide  with  water  : 

PI3  + 3H,0  = H3PO3  + 3 HI. 

A more  convenient  procedure  consists  in  adding  15  parts  of  iodine  (to  which  lo  parts 
of  water  have  been  added)  gradually  and  while  cooling  to  a mixture  of  one  part  of  red 
phosphorus  and  four  parts  of  water,  and  then  gently  heat  the  same  ; or  allow  an  emul- 
sion of  red  phosphorus  (5  parts)  with  water  (10  parts)  to  flow  gradually  and  at  first  very 
slowly  upon  iodine  (100  parts)  moistened  with  water  (lo  parts)  (Lothar  Meyer,  Ber.  20 
(1887),  3381).  Hydrogen  iodide  prepared  in  this  way  is  invariably  contaminated  with 
phosphorus  compounds.  The  pure  gas  can  only  be  made  by  the  method  mentioned 
above  : by  conducting  iodine  vapor  and  hydrogen  over  heated  platinum  sponge.  Water 
is  then  saturated  with  the  escaping  gas  and  on  heating  this  fuming  hydriodic  acid  a steady 
current  of  hydrogen  iodide  is  easily  obtained.  It  is  dried  by  passing  it  over  phosphorus 
pentoxide.  (See  Bodenstein,  Zeit.  f.  phys.  Ch.  13  (1894),  59.) 

Another  method  of  obtaining  aqueous  hydrogen  iodide  consists  in 
passing  hydrogen  sulphide  into  water  to  which  finely  pulverized  iodine 
is  added,  as  long  as  decolorization  occurs  : 

H^S  + 12  = 2HI  + S. 

Filter  off  the  separated  sulphur  and  distil  the  liquid.  Dry  iodine  and 
dry  hydrogen  sulphide  do  not  react  upon  one  another. 

Hydrogen  iodide  is  a colorless  gas;  it  fumes  strongly  in  the  air;  its 
density  is  128  (02  = 32)  or  4.4  (air=  i).  Under  a pressure  of  4 atmos- 
pheres (at  0°)  it  is  condensed  to  a liquid  which  boils  at  — 34°-*  It 
solidifies  at  lower  temperatures  and  remelts  at  — 51°.  It  is  easily  soluble 
in  water,  i volume  of  the  latter  dissolving  450  volumes  of  the  gas  at  10°. 
The  solution  saturated  at  0°  C.,  has  a specific  gravity  of  i 99,  and  fumes 
strongly  in  the  air.  If  the  solution  be  heated  hydrogen  iodide  is  ex- 
pelled, the  temperature  rises  and  at  126°  a solution  of  1.70  specific  gravity, 
containing  57  ])er  cent,  of  hydrogen  iodide,  distils  over.  Its  composi- 
tion corresponds  closely  to  the  formula  HI  5H2O. 

Hydrogen  iodide  is  a rather  unstable  compound.  Its  decomposition 
takes  place  at  all  temperatures  at  which  it  exists  as  a gas  ; the  speed  of  its 
disintegration  increases  rajiidly  with  the  temperature.  While  only  the 
two-thousandth  part  of  the  hydrogen  iodide  separates  into  hydrogen  and 
iodine  in  ninety  days  at  a temperature  of  100°,  almost  one-fourth  of  it 


* Consult  Estreicber,  Zeit.  f.  phys.  Cli.,  20  ('1896),  605,  upon  the  behavior  of  hydrogen 
chloride,  bromide  and  iodide  at  low  temperatures. 


64 


INORGANIC  CHKMISTRV. 


will  be  decomposed  in  fifteen  minutes  at  518°,  and  tlie  maximum  decom- 
position for  tins  temperature  will  then  have  been  attained  (see  Ibssoci- 
ation  of  Water).  At  high  temperatures  oxygen  decomposes  into  water 
and  iodine  : 

2lII  -t-  O^IbO  + Ij. 


On  bringing  a flame  near  the  mouth  of  a vessel  containing  a mixture 
of  hydrogen  iodide  and  oxygen,  violet  iodine  vapors  will  be  liberated. 
The  same  will  be  noticed  when  fuming  nitric  acid  is  drojijied  into  a vessel 
containing  the  gas;  in  this  reaction  the  oxygen  of  the  acid  oxidizes  the 
hydrogen  and  liberates  iodine.  The  nitric  acid  breaks  down  into  com- 
pounds containing  less  oxygen.  All  oxidizing  bodies  behave  in  the  same 
way;  the  hydrogen  iodide  abstracts  their  oxygen  and  reduces  The 

same  fact  is  noticed  with  concentrated  suliihuric  acid  when  theattemjit  is 
made  to  apply  it  in  liberating  hydrogen  iodide  from  an  iodide.  The 
oxygen  of  the  air  gradually  decomposes  aqueous  hydrogen  iodide  at  the 
ordinary  temperature,  and  es])ecially  in  sunlight.  The  solution,  at  first 
colorless,  becomes  brown,  owing  to  seiiaration  of  iodine,  which  in  the 
beginning  dissolves;  subsequently,  however,  it  separates  in  beautiful 
crystals. 

At  ordinary  temperatures  mercury  and  silver  decompose  hydrogen 
iodide,  with  separation  of  hydrogen: 

2lIl4-2Ag  = 2Agl4-H2. 

Chlorine  and  bromine  liberate  iodine  from  hydrogen  iodide  (see 
P-  54)- 

This  compound  is  employed  as  a powerful  reducing  agent  in  laboratory 
work. 


4.  HYDROGEN  FLUORIDE. 

HFl  = 20.01. 

It  is  obtained,  like  hydrogen  chloride,  by  decomposing  fluorides  with 
sulphuric  acid.  Finely  pulverized  fluors-par  (CaFb)  is  mixed  with  con- 
centrated sulphuric  acid  and  heated  gently  : 

CaFb  -f  H.,S04  = CaSO^  + 2HFI. 

Calcium  Calcium 

fluoride.  sulphate. 

The  oiieration  is  executed  in  a lead  or  platinum  retort,  as  the  hydrogen 
fluoride  attacks  glass  and  most  of  the  metals.  The  esca])ing  gas  is  con- 
ducted into  water.  To  get  perfectly  anhydrous  hydrogen  fluoride,  heat 
hydrogen  ])otassium  fluoride,  which  then  decomposes  according  to  the 
following  equatifin  : 

IIKFI,  = KFl  + 1 1 FI. 

Anhydrous  hydrogen  fluoride  is  a colorless,  very  mobile  liquid,  fuming 
strongly  in  the  air,  and  attracting  moisture  with  avidity;  it  boils  at 
-{-19.4°  (’.,  and  has  a specific  gravity  of  0.98  at  12°.  To  recondense 
llie  gas  it  must  be  cooled  to  — 20°.  Hydrogen  fluoride  solidifies  at 
— 102.5°  and  rcmells  at  — 92.5°. 


HYDROGEN  FLUORIDE. 


65 


The  gas  density  of  hydrogen  flouride  equals  20.01  (O,^  — 32)  at  100°,  corresponding  to 
the  molecular  formula  HFl.  At  30°,  however,  it  is  twice  as  large,  equaling  40.  It  fol- 
lows, therefore,  that  the  molecules  of  the  gas  at  the  latter  temperature  correspond  to  the 
formula  H2Fl2»  consist  of  two  chemical  molecules  of  HFl  (compare  Arsenic  Tri- 
oxide), 

The  concentrated  aqueous  solution  fumes  in  the  air ; when  heated, 
hydrogen  fluoride  escapes;  the  boiling  temperature  increases  regularly 
and  becomes  constant  at  120°,  when  a solution  distils  over,  the  specific 
gravity  of  which  is  1.15,  and  its  percentage  of  hydrogen  fluoride  is  35.3. 
'The  vapors  as  well  as  the  solution  are  poisonous,  extremely  corrosive, 
and  produce  painful  wounds  upon  the  skin. 

Hydrofluoric  acid  dissolves  all  the  metals,  excepting  lead,  gold  and 
platinum,  to  form  fluorides.  It  decomposes  all  oxides,  even  the  anhy- 
drides of  boric  and  silicic  acids,  which  it  dissolves  to  form  boron  and 
silicon  fluorides.  Glass,  a silicate,  is  also  acted  upon;  hence  the  use  of 
the  acid  for  etching  this  substance  (compare  Silicon  Fluoride).  To  do 
this,  coat  the  glass  with  a thin  layer  of  wax  or  paraffin,  draw  any  figure 
upon  it  with  a pin,  and  then  expose  it  to  the  action  of  the  gaseous  or 
liquid  hydrogen  fluoride.  The  exposed  portions  appear  etched  ; gaseous 
hydrogen  fluoride  furnishes  a dim,  and  liquid  hydrogen  fluoride  a smooth, 
transparent  etching. 

Vessels  of  lead,  platinum,  or  caoutchouc  are  employed  for  the  preser- 
vation of  hydrofluoric  acid,  as  they  are  not  affected  by  it. 


These  halogen  derivatives  of  hydrogen  show  great  resemblance  to  one 
another.  At  ordinary  temperatures  they  form  strongly  smelling  and 
fuming  gases,  which  can  be  easily  condensed  to  liquids.  Their  fuming 
in  moist  air  is  due  to  the  fact  that  they  are  dissolved  by  the  water 
vapor  and  the  resulting  solutions  appear  as  a cloud  consisting  of  very 
minute  drops.  Being  readily  soluble  in  water,  they  are  only  partly  ex- 
pelled from  their  saturated  solutions  by  boiling;  solutions  of  definite 
composition  distil  over,  but  these  cannot  be  regarded  as  definite  chemical 
combinations  of  the  halogen  hydrides  with  water,  because  their  com- 
position depends  upon  the  pressure  at  which  they  are  boiled  (Roscoe). 

As  acids  they  neutralize  the  bases  and  form  haloid  salts,  which  also  re- 
sult by  the  direct  union  of  the  halogens  with  metals. 

The  densities  of  the  halogen  hydrides  exhibit  a gradation  similar  to 
that  of  the  densities  of  the  halogens  (p.  56)  : 

HFl  HCl  HBr  HI 

Densities,  20.01  36.46  80.97  127.86 

The  difference  in  chemical  deportment  corresponds  to  this  gradation. 
Hydrogen  fluoride  is  the  most  stable,  and  acts  most  energetically  ; fluo- 
rine unites  in  the  dark  with  hydrogen  ; chlorine  combines  with  it  in  sun- 
light, while  bromine  and  iodine  require  higher  temperatures  for  their  com- 
bination with  it.  On  the  other  hand,  hydrogen  iodide  is  decomposed  at  a 
gentle  heat  (180°),  into  its  constituents;  the  more  stable  hydrogen  bro- 
mide at  800°,  while  hydrogen  chloride  remains  unaltered  up  to  1500°  C. 


66 


INORGANIC  CHEMISTRY. 


Corresponding  to  this  we  liave  the  very  energetic  action  of  fluorine,  and 
the  tolerably  ready  action  of  chlorine  iii)on  water,  oxygen  sej)arating  at 
the  same  time : 

n/)  -p  Cl,  ^ 2IICI  -f  O. 

Iodine  does  not  act  upon  water.  The  opjjosite  reaction  occurs  : oxygen 
decomi)oses  hydrogen  iodide  into  water  and  iodine: 

2lII  -f  O = II/)  + I,. 

Bromine  occupies  an  intermediate  position  between  chlorine  and 
iodine;  in  dilute  acpieous  solution  it  decomposes  water  into  hydrogen 
bromide  and  oxygen,  while  a concentrated  solution  of  hydrogen  bromide, 
on  the  contrary,  is  partly  decomposed  by  oxygen  into  water  and  free 
bromine. 

From  all  theabove  it  is  evident  that  the  affinity  of  fluorine  for  hydrogen 
is  the  greatest;  then  follow  chlorine  and  bromine,  and  finally,  as  the 
least  energetic  element,  we  have  iodine  (see  j).  57).  Fluorine  holds 
an  exceptional  position  with  the  other  halogens  in  that  its  hydride  is  a 
liquid  at  the  ordinary  temperature,  and  many  of  its  metallic  derivatives 
show  a solubility  directly  oiiposite  to  that  of  the  metallic  chlorides,  bro- 
mides, and  iodides.  This  will  be  discussed  later. 


THERMO-CHEMICAL  DEPORTMENT  OF  THE  HALOGENS. 

The  quantities  of  heat,  disengaged  or  absorbed  in  chemical  reactions, 
afford  the  most  satisfactory  explanations  of  the  deportment  of  the 
halogens  with  hydrogen,  and  indeed  of  all  the  chemical  elements  and 
compounds  toward  one  another.  These  heat  changes  are  also  called 
positive  and  negative  thermal  values  (Jieat  fnodulus')  (see  p.  30). 

The  quantities  of  heat  are  estimated  in  heat  units  or  calories.  The  quantity  of  heat 
required  to  raise  one  gram  of  water  1°  C.  from  15°  (measured  with  an  air-thermometer),  is 
taken  as  the  heat  unit  (small  calorie),  or  a thousand  times  this  quantity  can  be  taken  ; 
then  it  would  be  the  quantity  of  heat  needed  to  raise  i kilogram  of  water  from  15°  to  16° 
(large  calorie.  Cal.).  Large  calories  will  be  used  in  the  following  pages.  (See  further, 
Nernst,  Theoretische  Chemie,  2 Aufl.  (1898),  10.) 

To  obtain  data  that  may  be  easily  compared,  the  quantities  of  heat  are 
referred  to  qtiantities  in  grams  corresponding  to  the  atomic  or  molecular 
weights  of  the  elements  entering  into  combination.  Thus,  in  the  union 
of  19  grams  of  fltiorine  (FI  19)  with  i.oi  grams  of  hydrogen  (H  = i.oi) 
to  form  20.01  grams  of  hydrogen  fluoride  (HFl  = 20.01),  37.6  Cal.  are  set 
free,  and  in  the  formation  of  36.46  grams  of  hydrogen  chloride  (HCl  = 
36.46)  from  its  elements  22.0  Cal.  When  79.96  grams  of  bromine 
( Br  = 79.96)  combine  with  i.oi  grams  of  hydrogen  to  form  hydrogen 
bromide  8.4  Cal.  are  developed,  while  in  the  formation  of  127.86  grams 
of  hydrogen  iodide  from  solid  iodine  and  hydrogen  6.0  Cal.  are  ab- 
sorbed, but  with  iodine  vajior  and  hydrogen  only  1.5  Cal.  (4.5  Cal.  being 
required  for  the  vaporization  of  the  given  amount  of  iodine). 


THERMO-CHEMICAL  DEPORTMENT  OF  THE  HALOGENS. 


67 


This  may  be  expressed  according  to  the  method  of  J.  Thomsen,  as  follows  : 

(H,F1)  = -f37.6  Cal.  ; (H,C1)  == -f  22.0  Cal. ; (H,Br)  = +8.4  Cal. ; 

(H,I)=-i.5Cal. 

The  first  three  reactions,  in  which  heat  is  liberated,  are  exotherDiic^  while  the  heat- 
absorbing combination  of  iodine  with  hydrogen  represents  an  endother77iic  reaction  (see 
p.  30).  The  energy-content  of  hydrogen  fluoride,  chloride  and  bromide  is  less,  and  that 
of  hydrogen  iodide  greater  than^  that  of  their  components. 

The  quantity  of  heat  disengaged  in  a combination  must  not  be  regarded  as  a measure 
of  the  chemical  affinity.  As  the  elements  with  few  exceptions  ; see  argon,  helium,  mer- 
cury, and  cadmium — do  not  exist  as  free  atoms,  but  as  molecules  these  require  a definite 
quantity  of  heat  to  decompose  them  into  atoms  before  they  can  enter  into  chemical  reac- 
tion. This  necessitates  a definite  amount  of  work  (addition  of  energy).  The  union  of 
chlorine  with  hydrogen  proceeds  according  to  the  molecular  equation  (p.  76)  : 

HH-f  C1C1  = 2HC1o 

The  heat  here  disengaged  (2  X 22.0  =r  44.0  Cal.)  is  the  algebraic  sum  of  the  follow- 
ing unknown  thermal  values  : (i)  — x Cal.,  required  for  the  decomposition  of  the  hydro- 
gen molecule  into  free  atoms  ; (2)  — y Cal.,  consumed  in  the  decomposition  of  the  chlorine 
molecule;  (3)  -fz  Cal.,  liberated  in  the  formation  of  hydrogen  chloride  from  the  free 
chlorine  and  hydrogen  atoms  ; hence  z — x — y = 44.0,  i.  e. , the  thermal  value  is  deduced 
from  the  three  unknown  values. 

It  is  very  probable  that  the  union  of  the  free  atoms  always  occurs  with  heat-disengage- 
ment, and  the  heat-absorption,  observed  in  many  chemical  changes,  is  to  be  credited  to 
the  decompositions  to  which  attention  has  been  directed.  In  the  formation  of  hydrogen 
iodide,  for  example,  the  sum  of  — x and  — y is  very  probably  greater  than  z. 

The  greater  the  heat  developed  in  a reaction,  the  more  energetically 
and  the  more  readily  will  it  occur,  and  in  general,  the  resulting  com- 
pounds will  be  the  more  stable  (compare  p.  65,  Behavior  of  the 
Halogen  Hydrides).  The  energetic  reactions,  those  which  are  accompa- 
nied by  very  appreciable  evolution  of  heat,  must  be  viewed  as  transitions 
of  systems  from  a state  of  comparative  instability  to  one  of  greater  per- 
manency. The  opposite  occurs  if  the  chemical  change  only  takes  place 
upon  the  addition  of  external  energy.  The  compound  then  produced 
passes  very  readily  into  the  original  and  more  stable  system.  In  this 
sense  the  principle  of  greatest  heat-development  is  true  (p.  31). 

In  this  way  it  can  be  understood  from  the  thermal  values  of  the  halogen 
hydrides  why  it  is  that  iodine  is  displaced  by  the  other  halogens,  bromine 
by  chlorine  and  fluorine  and  chlorine  by  fluorine  from  their  hydrides  and 
their  metallic  derivatives — corresponding  to  the  following  thermo-chem- 
ical equations : 

HI  + Cl  =r  HCl  + I . . . (+28  Cal.) 

( — 6.0)  (22.0) 

IIBr-|-Cl=:HCl  + Br  . . . (-f  13.6  Cal.) 

(8.4)  (22.0) 

HCl  -f  Fl=  IIFl  + Cl  . . .(+15.6  Cal.) 

(22.0)  (37.6) 

The  thermo-chemical  sign  of  a reaction  is  obtained  by  deducting  from  the  heat  of 
formation  of  the  jjroducts  that  of  those  reacting. 

The  reactions  do  not  proceed  wholly  in  the  sense  indicated.  They  are  limited  by 
accurately  opposing  forces.  When  bromine  acts  upon  hydrochloric  acid  the  transposition 
takes  place,  if  even  in  a very  slight  degree,  according  to  the  equation  Br  -f-  HCl  = 


68 


INORCIANIC  CHEMISTRY. 


IIBr  -(-  Cl.  And  it  must  therefore  be  assumed  from  tliis  equation  that  when  chlorine 
acts  upon  liydrol)romic  acid  a portion  of  the  tnomine  liberated  will  trans])ose  itself  with 
the  hydrogen  chloride  which  has  been  formed  : the  reaction  reverses  itself,  is  retrogres- 
sive. To  what  extent  this  hai)pens,  to  what  degree  of  divisions  these  reactions,  occurring 
simultaneously  but  proceeding  in  opposite  directions,  will  go,  depends  uj)on  the  nature 
of  the  reacting  bodies,  the  ratios  of  their  quantities  in  unit  volume,  the  temperature,  the 
time,  the  pressure  ; and  finally  a state  of  ecpiilibrium  will  be  developed.  'I'lie  retrogres- 
sive reaction  very  frecjiiently  advances  so  slowly,  conse(|uently  amounts  to  so  little,  that 
it  may  ajjpear  as  if  it  actually  did  not  occur  at  all. 

These  facts  demonstrate  that  Terthelot’s  principle  of  the  greatest  heat  development  is 
not  universally  correct  although  it  often  indicates  the  direction  in  which  a transhnniatioji 
occurs  easily  and  completely.  As  an  argument  favoring  this  we  have  the  different  decorn- 
])osability  of  the  gaseous  halogen  hydrides  by  oxygen,  taking  into  consideration,  of  course, 
the  heat  of  formation  of  water.  For  acjueous  vapor  this  is  57.2  Cal.,  while  for  the  liquid 
it  is  68.3  Cal.  The  union  of  i gram  of  hydrogen  with  8 grams  of  oxygen  is  attended 

with  a thermal  value  of  — 28.6  Cal.,  which  is  greater  than  that  of  hydrogen  chloride, 

bromide  or  iodide  but  less  than  that  of  hydrogen  fluoride.  Consequently,  oxygen  should 
displace  chlorine,  bromine  and  iodine  but  not  fluorine  from  their  re.spective  hydrides,  and 
this  takes  place  the  more  readily  the  greater  the  difference  in  the  heat  of  formation. 

In  fact,  we  observed  (p.  64)  that  when  a flame,  or  some  glowing  substance,  was  brought 
in  contact  with  a mixture  of  hydrogen  iodide  and  oxygen,  all  the  iodine  was  separated  in 
the  form  of  vapor,  in  accordance  with  the  following  equation  : 

2111  + O = H.p  (vapor)  H 2I  . . . ( + 69.2  Cal.) 

Oxygen  also  liberates  bromine  from  hydrogen  bromide  at  a temperature  of  about  500° 
(neither  hydrogen  bromide  nor  water  suffer  di.ssociation  at  this  temperature).  Aqueous 
vapor  is  also  produced.  Hydrogen  chloride,  however,  is  only  partially  decomposed  by 
oxygen  even  at  higher  temperatures. 

Phirther,  in  accordance  with  this  idea,  in  a mixture  of  chlorine,  hydrogen  and  oxygen, 
the  hydrogen  will  first  unite  with  chlorine  and  if  any  remain  then  with  oxygen,  although 
the  heat  of  formation  of  water  (H2,0)  = 57.2  Cal.  is  greater  than  that  of  hydrogen 
chloride  (Cl2,H2')  =44.0  Cal.  For  a more  exhaustive  study  of  thermo-chemical  relations 
the  student  is  referred  to  H.  Jahn’s  Grundsatze  der  Thermochemie  (2  Aufl.  Wien,  1892), 
and  also  to  the  works  of  Ostwald  and  of  Nernst  (see  pp.  49,  66). 


COMPOUNDS  OF  THE  HALOGENS  WITH  ONE 
ANOTHER. 

These  compounds,  formed  by  the  union  of  the  halogens  with  one 
another,  are  very  unstable,  and  it  may  be  remarked  here,  that  this  is  also 
true  of  most  derivatives  obtained  from  elements  which  are  chemically 
similar. 

When  chlorine  is  conducted  over  dry  iodine,  the  latter  being  in  excess, 
iodine  monochloride  results,  and  when  the  chlorine  is  in  excess,  iodine 
trichloride  is  formed. 

Iodine  Monochloride — ICl — is  a red  crystalline  mass,  fusing  at  24.7°,  and  distilling 
a little  above  100°.  Water  decomposes  it  easily,  with  formation  of  iodic  acid,  iodine, 
atid  hydrogen  chlorid(;.  If  fused  iodine  chloride  be  allowed  to  .solidify  .slowly  at  a low 
temperature  { — I(j°)  a modification,  melting  at  -|  14°,  is  produced.  The  latter,  however, 
reverts  very  readily,  with  heat  evolution,  to  the  higlier  melting  body. 


WEIGHT  PROPORTIONS. 


69 


Iodine  Trichloride — 10)3 — is  formed  upon  mixing  iodic  acid  with  concentrated 
hydrochloric  acid,  and  by  the  action  of  jdiospliorus  pentachloride  upon  iodic  anhydride. 
It  crystallizes  in  long,  yellow  needles,  and,  when  heated,  suffers  decomposition  into  iodine 
chloride  and  chlorine  (at  ordinary  pressure,  the  dissociation  commences  at  25°).  It  dis- 
solves in  a little  water  without  alteration  ; but  large  quantities  cause  partial  decomposi- 
tion, with  formation  of  iodic  and  hydrochloric  acids. 

Iodine  Bromide — IBr — obtained  by  the  direct  union  of  the  elements,  consists  of 
iodine-like  crystals,  fusing  at  about  36°. 

Iodine  Pentafluoride — IFl^ — is  produced  by  the  action  of  iodine  upon  silver  fluoride, 
and  forms  a colorless,  strongly  fuming  liquid. 


WEIGHT  PROPORTIONS  IN  THE  UNION  OF  THE  ELEMENTS. 

ATOMIC  HYPOTHESIS.  CHOICE  OF  ATOMIC  WEIGHTS. 

In  the  preceding  pages  several  different  and  independent  methods  have 
been  described  for  the  preparation  of  each  of  the  halogen  hydrides.  But 
it  is  immaterial  which  of  these  may  be  selected  in  making  any  of  these 
compounds,  for  if  the  product  be  carefully  purified  and  then  analyzed  the 
hydrogen  and  the  halogen  will  always  be  present  in  a definite,  unalterable 
proportion  by  weight.  The  percentage  composition  of  the  pure  halogen 
hydrides  will  be  found  under  all  circumstances  to  be  the  following  : 

H 5.05  H 2.77  H 1.25  H 0.79 

FI  94.95  Cl  97.23  Br  98.75  I 99.21 

HFl  100.00  IICl  106.00  IIBr  100.00  HI  100.00 


A regularity  exists  here,  which  prevails  in  all  chemical  compounds. 
For  it  is  not  alone  in  the  halogen  hydrides,  but  without  exception  m 
every  chemical  combination  that  the  constituents  occur  in  definite  unalter- 
able proportions  by  weight.  This  observation  ascertained  by  experiment 
and  based  u])on  facts  has  been  called  the  law  of  definite  or  constant  pro- 
portions. From  the  period  in  which  the  French  chemist  Louis  Proust 
victoriously  defended  it  against  the  attack  of  his  countryman  Claude 
Louis  Berthollet,  the  great  theorist,  in  a remarkable  controversy  waged 
from  1799-1807,  down  to  the  present  no  fact  has  been  observed  which 
contradicts  the  law. 

It  will  be  advisable  for  later  considerations  that  the  quantities  of  the 
halogens  be  calculated  from  the  numbers  given  above  for  the  percentage 
composition  of  the  haloid  acids,  which  combine  with  a definite  quantity 
by  weight  of  hydrogen,  the  constituent  common  to  these  four  compounds. 
For  reasons  to  be  discussed  later  hydrogen  is  no  longer  taken  as  unit,  as 
has  been  done  in  the  past,  and  it  will  be  better,  therefore,  to  compare  the 
quantities  by  weight  of  the  elements  with  one  another,  which  are  capable 
of  uniting  with  i.oi  ])arts  by  weight  of  hydrogen.  In  this  manner  we 
arrive  at  the  following  numbers: 

II  I.OI  II  I.OI  II  I.OI  II  I.OI 

FI  19.  Cl  35.45  Br  79.96  I 126.85 


IICl  36.46  IIBr  80.97 


III  127.86 


1 1 FI  20.01 


70 


INORGANIC  CUKMISTRY. 


This  simple  recalciilalion  arfords  a dearer  insiglit  into  the  ratios  by 
weight  according  to  whicli  the  halogens  unite  with  hydrogen.  We  thus 
discover  that  19  parts  of  lluorine,  35.45  parts  of  chlorine,  79.96  jarts  of 
bromine,  and  126.85  P-'ii'ts  of  iodine,  inasmuch  as  they  are  capable  of 
uniting  with  i oi  parts  of  hydrogen,  are  eipial  or  e(pii valent  to  one 
another.  These  numbers  answer  not  only  for  the  derivatives  of  the  halo- 
gens with  hydrogen,  but  we  find  them  in  the  compounds  of  the  halogens 
with  one  another,  and  in  their  derivatives  with  other  elements,  'rims  in 
iodine  monochloride  the  chlorine  and  iodine  are  present  in  proportions 
by  weight  35.45  : 1 26.85  iodine  monobromide  the  ratio  of  bromine 

and  iodine  is  79.96:  126.85.  grams  of  sodium  be  converted  into 

fluoride,  chloride,  bromide,  and  iodide  the  (quantities  of  the  halogens 
required  for  this  purpose  will  again  be  in  the  ratio  of  19  : 35.45  : 79.96: 
126.85.  same  occurs  if  we  substitute  potassium,  calcium,  magne- 

sium, zinc,  silver,  etc.,  for  sodium.  Thus  19  q)arts  by  weight  of  fluorine 
combine  with  the  following  weights  of  the  metals:  23.05  i)arts  sodium, 
39.15  quarts  potassium,  32.7  quarts  zinc,  31.8  quarts  copper,  100. 15  q^iarts  mer- 
cury,— and  35.45  parts  chlorine,  79.96  q)arts  bromine,  and  126.85 
iodine  combine  with  exactly  the  same  quantities  by  weight  of  these 
metals. 

Let  us  take  another  example.  On  bringing  copper  into  a solution  of  a 
mercuric  salt  the  former  dissolves,  while  mercury  seq)arates  out;  indeed, 
31.8  parts  of  copq^er  displace  100. 15  q:)arts  of  mercury.  If  zinc  be  brought 
into  the  copper  solution  thus  obtained,  it  will  dissolve,  while  coqjper 
seq^arates — and  32.7  quarts  of  zinc  seq^arate  31.8  quarts  of  copper.  Further- 
more, zinc  disqdaces  the  hydrogen  in  acids;  from  all  of  them  32.7  parts 
of  zinc  seq^arate  i.oi  quarts  of  hydrogen.  In  all  these  reactions  we  ob- 
serve the  elements  apq^earing  in  the  same  quantities  by  weight. 

There  is  a net  of  q^erfectly  definite  proportions,  by  weight,  connecting 
all  these  bodies  with  one  another,  and  also  the  reactions  which  occur 
between  them,  i.oi  parts  of  hydrogen  combine  with  35.45  qiarts  of 
chlorine,  and  this  quantity  of  the  latter  with  23.05  parts  of  sodium,  31.8 
parts  of  copper,  32.7  parts  of  zinc.  32.7  qiarts  of  the  latter  rnetal  precipi- 
tate 31.8  parts  of  copper  and  100.15  parts  of  mercury,  from  their  salt  solu- 
tions, and  these  quantities  of  the  two  metals  are  caqiable  of  uniting  with 
the  quantities  of  fluorine,  bromine,  and  iodine  which,  like  the  35.45  parts 
of  chlorine,  combine  with  23.05  qiartsof  sodium  or  i.oi  parts  of  hydrogen, 
etc.  Proceeding  in  this  way  with  i.oi  qiarts  by  weight  of  hydrogen  we 
obtain  a number  for  each  element  which  may  be  called  its  co7?ibinin!:!; 
7ueight.  It  will  be  discovered  that  the  atomic  weight  of  oxygen  (O  = 16) 
is  the  basis  of  all  these  number  ratios.  What  has  been  written  may  be 
summarized  thus:  The  elements  combine  with  one  anothe)'  in  the  ratio  of 
their  combining  weights. 

A series  of  very  im])ortant  facts,  however,  compels  us  to  accord  this  law 
a broader  meaning.  It  often  transq)ircs  that  two  elements  combine  with 
one  another  not  only  in  one  pro|)ortion  by  weight,  as  in  the  case  of  the 
halogens  and  hydrogen,  but  in  several  ratios,  'rims,  there  are  two  com- 
I)ounds  of  chlorine  with  iodine — the  monochloride  and  the  trichloride. 
The  first  of  these  always  contains  35.45  qxarts  by  weight  of  chlorine  to 


WEIGHT  PROPORTIONS. 


71 


126.85  parts  by  weight  of  iodine,  and  the  second  3 X 35-45  = 106.35 
parts  of  chlorine  to  the  same  quantity  of  iodine.  Two  compounds  of 
hydrogen  and  oxygen  are  known  : water  and  hydrogen  peroxide.  In 
water  there  are  always  8.00  parts  by  weight  of  oxygen  to  i.oi  parts  by 
weight  of  hydrogen,  while  in  the  peroxide  to  i.oi  parts  of  hydrogen 
there  are  8.00  X 2 = 16  parts  of  oxygen.  35.45  parts  of  chlorine  com- 
bine not  only  with  31.8  parts  of  copper  and  100.15  parts  of  mercury,  but 
also  with  63.6  parts  of  copper  and  200.3  parts  of  mercury.  Oxygen 
forms  five  distinct  compounds  with  nitrogen  with  the  following  propor- 
tions by  weight,  in  which  8,  the  combining  weight  of  oxygen  (see  above), 
is  made  the  basis  of  comparison  : 


Nitrogen.  Oxygen. 

Nitrous  Oxide, 14.04  parts  8 parts 

Nitric  Oxide, 14.04  “ 16  “ =2X8.00 

Nitrous  Anhydride, 14.04  “ 24  “ =3X8.00 

Nitrogen  Dioxide, 14.04  “ 32  “ =4X8.00 

Nitric  Anhydride, 14.04  “ 40  “ = 5 X 8.00 


Proust  observed  that  two  elements  could  combine  with  each  other  in 
different  proportions  by  weight  and  that  in  so  doing  the  composition 
changed  by  definite  increments.  The  underlying  law,  however,  was  first 
recognized  and  propounded  (evidently  as  the  result  of  atomic  considera- 
tions) by  John  Dalton,  and  definitely  established  on  a scientific  basis 
through  the  labors  of  J.  J.  Berzelius.  It  is  the  law  of  multiple  propor- 
tions : When  two  elements  unite  in  several  p7'oportio7is,  the  quantities  of 
the  seco7td  ele77te7it  co77ibined  with  definite  a77iounts  of  the  first  bear  a si77iple 
ratio7ial  ratio  to  each  other.  The  law  of  definite  combination  by  weight 
can  be  so  expanded  that  it  will  at  the  same  time  include  the  law  of  con- 
stant and  also  that  of  multiple  proportions.  Then  it  would  read:  The 
ele77ients  07ily  unite  m the  ratio  of  their  co77ibming  weights  or  simple  ratio7ial 
77iultiples  of  the  sa77ie. 

Compounds,  therefore,  contain  their  constituents  either  in  the  ratio  of 
their  combining  weights,  or  some  simple  rational  multiple  of  these  com- 
bining weights — a simple,  self-evident,  analytical  view  of  the  law. 

The  elements  of  the  doctrine  of  weight-combinations  may  be  observed 
and  gathered  from  the  investigations  of  the  German  chemists — Karl 
Friedrich  Wenzel  and  Jeremias  Benjamin  Richter  upon  the  neutralization 
of  bases  and  acids,  and  the  alternating  transposition  of  salts.  Following 
the  example  of  Richter  this  division  of  our  science  is  even  yet  designated 
stoichiometry  (r«  ffror/sJa,  the  constituents;  [xirpov,  measure),  and  the 
laws  just  deduced  are  the  stoichio77ietric  laws.  It  was  in  this  particular 
direction  that  Berzelius,  from  1808  forward,  achieved  so  much  by  many 
hundred  accurate  analyses  which  aided  very  materially  in  establishing 
a foundation  of  irreproachable  facts  for  the  jireceding  doctrine. 

[Wenzel:  Die  T.elire  von  der  Verwandtschaft,  Dre.sden,  1777  ; Richter:  Ueber  die 
neueren  (jegen.stande  der  Cheniie  ; lie.sonder.s  iin  7,  8,  and  9 Stuck,  1796-1798.  Com- 
pare also:  Rerzeliii.s’  Lehrbiich  der  (Jhemie,  5 Aufl.,  Bd.  iii,  1147;  and  Versuch,  die 
bestimmten  und  einfachen  Verhaltnisse  aufzufinden,  etc.,  in  Ostwald’s:  Klassiker  der 
exakten  Wi.ssenschaften,  Nr.  35.] 


72 


INORGANIC  CIIRMIS’l'kV. 


The  statements  tlnis  far  made  liave  l)een  free  from  assumption  and  liave 
been  i)roved  l)y  exi)eiiment  and  analysis;  they  find  malhemalieal  ex- 
pression in  tlie  law  of  chemical  pro])orti(ms.  JUit  now  speculation  enters. 
To  ex[)lain  the  remarkable  regulariiies  in  the  ratios  John  Dalton  took 
refuge  in  the  atomic  hypothesis — “one  of  the  greatest  steps  of  which 
chemistry  availed  itself  in  advancing  to  jierfection  ” — berzelius. 

As  previously  indicated  in  the  introduction  ([i.  25)  Dalton  assumed 
that  the  elementary  atoms  iireferred  uniting  in  the  ratio  of  1:1; 
and  whenever  but  one  compound  of  two  elements  ajipeared  Dalton  re- 
garded it  as  comiiosed  of  an  atom  of  each  of  the  two  elements.  If  several 
com])oiinds  existed  he  viewed  the  first  as  consisting  of  A -f-  T,  the  second 
of  A -|-  2B,  the  third  of  2 A -J-  B,  and  the  fourth  of  A -j-  3B  (New  Sys- 
tem of  Chemical  Philosophy,  vol.  1(1808)).  By  this  assumption  the 
facts,  expressed  in  the  stoichiometric  laws,  meet  with  an  astonishingly 
simple  explanation.  If  hydrogen  chloride  contains  one  atom  of  hydrogen 
to  one  atom  of  chlorine  its  composition  must  always  be  the  same:  law  of 
constant  proportions.  If  an  atom  each  of  hydrogen,  sodium,  potassium, 
silver,  etc.,  unite  with  an  atom  of  fluorine,  chlorine,  bromine,  and  iodine 
to  form  the  corres[)onding  fluoride,  chloride,  bromide,  and  iodide,  then 
in  all  these  compounds  the  fluorine,  chlorine,  bromine,  and  iodine  must 
appear  in  the  same  constant  combination  ratios,  i.  e.,  the  quantities  of 
fluorine,  chlorine,  bromine,  and  iodine,  which  unite  with  a definite  quan- 
tity of  another  element,  must  always  be  to  one  another  as  their  atomic 
weights:  law  of  combining  weights.  If  iodine  at  one  time  unites  with 
one  atom  and  again  wdth  three  atoms  of  chlorine,  the  quantities  of  the 
latter  which  iodine  in  one  instance  requires  to  yield  the  monochloride 
and  at  another  time  the  trichloride  stand  in  the  projiortion  of  i : 3 (Cl  : 
3CI):  law  of  7?iultiple  proportions. 

If  two  elements  combine  in  but  one  proportion  by  weight  then  the  re- 
sulting compound,  on  the  assumption  of  Dalton,  contains  one  atom  of  each 
element.  This  at  once  makes  it  possible  to  determine  the  ratio  of  the 
atomic  weights  of  these  elements.  Thus,  in  hydrogen  chloride  there  are 
97.23  percent,  of  chlorine  to  2.77  per  cent,  of  hydrogen,  and  on  the  as- 
sumption of  Dalton  the  compound  consists  of  one  atom  of  hydrogen  and 
one  atom  of  chlorine,  then  these  numbers  possess  a deejier  meaning,  for  in 
that  case  the  atomic  weights  of  the  two  elements  will  be  as  2.77  : 97.23. 
Similarly,  the  ratios  of  the  atomic  weights  of  the  other  elements  may  be 
determined  and  the  relative  ato7fiic  weights  would  re.sult  just  as  soon  as 
some  number  is  selected  as  the  atomic  weight  for  any  one  element  and 
the  ])roportion  numbers  are  referred  to  this  number.  The  comparison 
element  should  be  one  which  forms  analyzable  derivatives  with  most  of 
the  other  elements.  The  number  selected  for  its  atomic  weight  is  a 
matter  of  ])racticability  and  of  general  agreement.  These  are  the  reasons 
which  were  potent  in  the  choice  of  oxygen  as  the  standard.  Its  atomic 
weight  has  been  ])laced  at  t6.  Our  relative  atomic  weights,  therefore,  refer 
— Compare  Bericht  der  Kommission  fiir  die  Festsetzung  der 
Atomgewichte  (Landolt,  Ostwald,  Seul)ert),  Ber.  31  (1898),  2761. 

Hydrogen,  the  spocificnlly  liglilcst  cleiiuMif,  was  chosen  by  Dalton  as  the  standard  ; its 
atomic  weij^ht  was  taken  as  unit.  'I'lie  disadvantage  in  this  instance  is  tliat  hydrogen 


GENERAL  PROPERTIES  OF  GASES. 


73 


forms  comparable  derivatives,  allowing  of  accurate  analysis,  with  few  elements.  Conse- 
quently the  ratio  of  its  atomic  weight  to  that  of  other  elements  is  usually  determined  with 
the  aid  of  oxygen.  Since,  however,  the  ratio  of  oxygen  to  hydrogen  cannot  be  as  sharply 
determined  as  that  of  oxygen  to  the  atomic  weights  of  many  other  elements,  errors  which 
are  not  justihable  creep  into  the  atomic  weights  by  virtue  of  this  recalculation.  These 
reasons  led  Berzelius,  to  whom  we  aie  indebted  for  the  first  accurate  atomic  weight  deter- 
minations, to  reject  the  Dalton  unit.  He  chose,  with  analytical  acuteness,  oxygen  as  the 
standard,  setting  its  atomic  weight  at  lOO  (O  = looj. 

About  the  middle  of  the  present  century  a return  to  the  Dalton  hydrogen  unit  occurred. 
The  movement  was  inaugurated  mainly  by  the  French,  in  order  to  distinguish  new  theo- 
retical views  at  first  sight  from  the  older  notions  held  by  Berzelius  and  his  students.  The 
standard  has  met  with  almost  universal  adoption  from  that  period  down  to  the  present  time. 

In  a recalculation  of  the  atomic  weights,  conducted  by  L.  Meyer  and  Carl  Seubert 
(1893)  with  all  previous  determinations  as  basis,  the  atomic  ratio  O : H =:  15.96  : i was 
assumed  as  the  most  probable.  Hitherto  it  has  been  regarded  as  the  most  sati.sfactory  by 
the  majority  of  chemists  and  has  been  adopted  in  this  and  other  text-books.  Since,  how- 
ever, J.  Thomsen  (1895)  and  particularly  E.  W.  Morley  by  a series  of  most  admirable  and 
painstaking  experiments  have  redetermined  this  ratio,  the  unsatisfactoriness  of  the  hydro- 
gen unit  has  become  more  apparent.  All  atomic  weights  determined  in  ratio  to  oxygen 
are  affected  by  the  change  in  this  proportion  O : H.  To  avoid  similar  difficulties  in  the 
future — for  the  ratio  O : H cannot  be  regarded  as  definitely  decided — chemists  have  taken 
oxygen  as  the  element  of  comparison,  and  in  order  that  the  new  atomic  numbers  may  be 
as  similar  as  possible  to  those  previously  used,  the  suggestion  of  Ostwald  that  the  atomic 
weight  of  oxygen  be  placed  at  16  has  been  followed.  The  atomic  weight  of  hydrogen 
will  then  be  H = 1.008,  or  for  all  practical  purposes  H — i.oi.  Compare  the  historical 
data  in  the  Zeit.  f.  anorg.  Chem.  14  (1897),  250,  256,  and  also  the  report  of  the  Atomic 
Weight  Commission  to  which  reference  has  already  been  made. 

The  relative  atomic  weights,  determined  in  the  manner  indicated,  are, 
of  course,  only  correct  if  our  assumption  as  to  the  number  of  atoms  in 
union  with  one  another  is  correct.  The  chemical  analysis  of  water  shows 
it  to  consist  in  one  hundred  parts  of  11.2  parts  of  hydrogen,  and  88.8 
parts  of  oxygen,  hence  2.02  parts  of  hydrogen  correspond  to  16  parts  of 
oxygen.  If,  then,  water  contains  one  atom  of  oxygen  in  union  with  one 
atom  of  hydrogen  the  atomic  weight  of  hydrogen  would  be  2.02  (O  = 16). 
It  may  be  possible,  however,  that  water  consists  of  two  atoms  of  hydrogen 
and  one  atom  of  oxygen  or  of  one  atom  of  hydrogen  and  two  atoms  of 
oxygen,  etc.  In  the  first  case  the  atomic  weight  of  hydrogen  would  be 
I.OI  and  in  the  second  4.04,  etc.,  etc. 

Analytical  results  afford  nothing  positive  for  the  solution  of  this  diffi- 
culty. This  was  the  condition  in  which  the  question  relating  to  the 
magnitude  of  the  atomic  weights  existed  fifty  years  ago.  To  establish 
these  correctly,  various  views  were  allowed  to  prevail,  none,  however, 
with  positive  foundation.  The  question  can  only  be  solved  upon  a basis 
of  special  views  of  the  gaseous  condition  and  new  facts  lying  chiefly  in 
the  domain  of  organic  chemistry. 

Nothing  definite  is  known  as  to  the  actual  magnitude  of  the  atomic  weights.  This  has 
thus  far  been  an  unimportant  matter  for  purely  chemical  consideration. 

GENERAL  PROPERTIES  OF  GASES.  ATOMIC-MOLECULAR 

THEORY. 

Gases,  i.  <?.,  aeriform  bodies,  above  their  critical  temperature,  posse.ss 
properties  in  pronounced  degree  which  are  inde])endent  of  their  chem- 
ical composition.  As  regards  alterations  of  pressure  and  Icnqicrature  a 

7 


74 


inor(;anic  chemistry. 


given  volume,  for  examiile,  one  liter,  of  hydrogen  deports  itself  in  the 
same  regular  manner  as  one  liter  of  any  other  gas  under  tlie  same  condi- 
tions. Here  chemical  composition  no  longer  plays  a role.  Just  as  soon 
as  a substance  i)asses  into  the  gas  condition  it  acquires  the  proiierties 
expressed  by  the  laws  about  to  be  discussed. 

The  relations  between  a volume  of  gas  and  the  pressure  under  which  it 
exists  are  expressed,  for  like  temperatures,  in  a law  named  after  its  first 
discoverer,  Hoyle  (1662),  although  it  often  aiijiears  under  the  name  of 
Mariotte  (1679).  reads:  the  volume  occupied  tfy  a given  mass  of  gas 
under  different  pressures  is  inversely  as  the  pressure.  If  the  jiressure  is 
trebled  the  volume  will  be  diminished  to  one-third;  if  the  jiressure  falls 
to  one-half,  the  volume  will  double,  etc.  If  the  gas  volumes  be  re|)re- 
sented  by  v,  v'  and  the  pressures  by  ]>,  p',  the  law  can  be  algebraically 
expressed  in  the  equation 


or 


p : p'  = \ V 

p : V = p'  : v''  = Constant. 


That  is,  the  product  of  the  pressure  and  volume  of  a given  quantity  of 
gas,  with  constant  temperature,  is  unalterable. 

The  second  law  was  discovered  in  1802,  almost  simultaneously,  by  Gay- 
Lussac  and  by  Dalton.  It  reads:  the  volume  of  any  gas  increases 
or  0.003665  of  its  volume  at  0°,  for  every  degree  centigrade  which  its 
temperature  rises  under  the  same  pressure.  This  number,  the  coefficient 
of  expansion  of  gases,  is  ordinarily  represented  by  the  letter  a.  If  the 
volume  Vo  of  the  gas  at  0°  was  equal  to  i,  then  at  1°  it  is  equal  to 
1.003665,  at  273°  equal  to  2,  and  universally  at  t" 


Vt  = Vo  Vo  . 0.003665  . t = Vo  (l  + at). 


By  combining  the  two  laws  there  results  the  important  equation 
P V = po  Vo  (i  + at), 

indicating  the  relations  between  the  volume  v,  which  a gas  under  pressure 
p occupies  at  t“and  the  volume  v^  at  0°  under  the  pressure  p^.  We  shall 
return  to  this  law  when  the  measurement  of  gases  and  their  variation  from 
the  Boyle-Gay-Lussac  law  are  considered. 

A third  law  relates  to  the  volume  ratios  according  to  which  gases 
combine  chemically  with  one  another.  These  were  first  investigated  by 
Gay-Lussac,  partly  in  conjunction  with  A.  v.  Humboldt.  Gay-Lussac 
announced  his  observations  as  follows:  The  gases  combine  according  to 
simple  volume  ratios;  the  volume  of  a compound  gas  bears  a simple  ratio 
to  the  volume  of  its  constituents.  In  the  union  of  hydrogen  and  a halo- 
gen to  form  a halogen  hydride  it  was  found  that  one  volume  of  hydrogen 
and  one  volume  of  halogen  combine  to  two  volumes  of  the  halogen 
hydride.  Hence,  in  this  instance,  the  union  took  place  without  change 
of  volume;  the  halogen  hydride  occupied  the  space  previously  occupied 
by  its  constituents:  an  important  fact  for  the  chemical  atomic  theory. 
It  is  concluded  from  this  and  the  previously  mentioned  facts,  according  to 
which  liydrogen  and  halogen  unite  with  one  another  in  the  ratio  of  their 
combining  weights,  that  the  ratio  between  the  weights  of  eipial  volumes 


GENERAL  PROPERTIES  OF  GASES. 


75 


of  these  gases,  i.  e.,  of  their  densities,  must  be  the  same  as  that  between 
the  combining  weights.  Each  halogen  has  but  one  derivative  with  hydro- 
gen. Hence,  according  to  Dalton,  it  may  be  assumed  that  in  the  case 
of  the  halogens  the  combining  weight  and  atomic  weight  are  the  same. 
Then  their  gas  density  would  be  proportional  to  their  atomic  weights. 
Here,  again,  and  also  in  similar  deportments  of  the  simple  as  well  as 
the  compound  gases,  expressed  in  the  law  relating  to  gases,  we  are  forced 
to  the  assumption  that  in  equal  volumes  of  different  gases  there  is  an  equal 
number  of  atoms — smallest  particles.  Dalton  and  Berzelius  both  assumed 
this  conception,  but  soon  abandoned  it  because  it  did  not  lead  to  a defi- 
nite, well-defined  distinction,  justified  by  facts,  between  elementary 
and  compound  gases.  Equal  volumes  of  hydrogen  and  of  chlorine  unite 
without  any  change  in  volume  to  hydrogen  chloride.  If  looo  atoms  are 
present  in  a given  volume  of  hydrogen,  looo  atoms  of  chlorine  would  be 
required  in  an  equal  volume,  and  by  their  union  there  would  result  looo 
particles  of  hydrogen  chloride,  which  would  occupy,  as  demonstrated  by 
experiment,  the  same  volume  as  the  chlorine  and  hydrogen  together : 

: 

looo  H -f-  looo  Cl  = looo  HCl. 

I vol.  I vol.  2 vols. 

A volume  of  hydrogen  chloride,  therefore,  would  contain  only  half  as 
many  particles  as  a like  volume  of  the  elementary  gases — a conclusion 
which  contradicts  everything  known  relative  to  the  constant,  physical 
deportment  of  the  gases  (being  independent  of  chemical  constitution) 
with  reference  to  the  laws  of  Boyle  and  Gay-Lussac. 

To  solve  this  contradiction  has  required  almost  a half  century  of  chem- 
ical investigation.  It  was  especially  the  advances  in  the  domain  of 
organic  chemistry  which  led  to  the  assumption  that  the  elementary  sub- 
stances as  well  as  the  compound  did  not  consist  of  a mass  of  free  atoms 
but  of  an  aggregation  of  atom-groups — of  molecules  (p.  24).  This  had 
been  advocated  (enunciated)  by  Avogadro  (1811),  then  by  Ampere 
(1814),  and  subsequently  by  Dumas  (1829),  but  it  was  only  at  the  close 
of  the  ’50’s  that  the  theory,  after  being  supported  by  the  Parisian  chem- 
ist Gerhardt  on  a purely  chemical  basis,  received  general  acceptance  and 
favor  [Gerhardt,  Lehrbuch  der  organ.  Chemie,  deutsche  Ausgabe,  4 
('1857),  598,  627,  etc.].  It  is,  consequently,  necessary  to  distinguish 
between  atom  and  molecule  {jjiolecula,  mass-particles).  It  is  not  the 
atoms  but  the  molecules  which,  as  a rule,  exist  in  the  gases  as  the  smallest 
constituent  particles  (exceptions  will  be  discovered  later;  see  also  p.  55 
on  the  Dissociation  of  the  Halogens).  It  is  obvious  that  the  smallest  par- 
ticles of  a compound  body  consist  of  several  atoms.  By  further  subdivi- 
sion such  a molecule  breaks  down  into  dissimilar  constituents:  a molecule 
of  hydrogen  chloride  is  resolved  into  hydrogen  and  chlorine.  Even  the 
elements,  when  free,  consist,  as  a rule,  of  molecules,  and  the  latter  of  sev- 
eral, generally  of  two,  atoms.  Hence,  the  previously  deduced  rule  that 
in  equal  volumes  of  the  elementary  gases  there  is  contained  an  equal  num- 
ber of  atoms  must  be  changed  and  amplified  somewhat  as  follows: 

The  number  of  molecules  in  a unit  volume  of  all  gases  is  the  same — 
like  pressure  and  like  temperature  being  presupposed.  Whde  not  entirely 
correct,  this  law  is  very  frequently  designated  as  the  law  of  Avogadro. 


76 


INORGANIC  CHEMISTRY. 


The  contradiction  between  the  conclusions  deduced  from  cliemistry 
and  from  physics  now  vanishes. 

'The  [)rocess  of  tlie  combination  of  hydrogen  with  chlorine  (and  the 
other  halogens)  must  be  conceived,  therefore,  to  be  somewhat  like  the 
following:  one  molecule  of  hydrogen,  containing  at  least  two  atoms  of 
hydrogen,  acts  ui)on  one  molecule  of  chlorine,  also  composed  of  at  least  two 
atoms  of  chlorine,  and  there  result  two  molecules  of  hydrogen  chloride  : 

II2  ( Cl,  — 2IICI. 

i.  e.,  hydrogen  chloride  contains  just  as  many  molecules  in  an  equal  vol- 
ume as  hydrogen  and  chlorine.  This  is  apparent  from  the  folh)wing 
representation,  which  attaches  itself  to  the  example  given  on  j).  75  : 


500  H, 


I volume. 


+ 


500  Cl, 


I volume. 


500  IICl 


500  IICl 


2 volumes. 


In  a similar  manner  two  volumes  of  hydrogen  (containing  2n  molecules) 
give  with  one  volume  of  oxygen  (containing  n molecules)  two  volumes  of 
aqueous  va])or ; consequently,  2fi  molecules  of  water.  In  2>^  molecules  of  the 
latter  (H^O)  there  are  contained  at  least  271  atoms  of  oxygen  ; therefore, 
in  71  molecules  of  oxygen,  271  atoms  of  oxygen — or  one  oxygen  molecule 
co7isists  of  at  least  two  ato77is  (compare  pp.  73,  80). 


2 volumes. 


nO, 

yield 

nH,0 

nH20 

I volume.  2 volumes. 


Let  us  take  another  analytical  example.  Two  volumes  (^2n  mol.)  of 
ammonia  gas  break  down,  under  the  influence  of  the  electric  spark,  into 
one  volume  ft  mol.)  of  nitrogen  and  three  volumes  of  hydrogen  (each  7t 
mol.).  From  this  it  is  evident,  as  in  the  preceding  example,  that  the 
nitrogen  molecule  also  is  composed  of  at  least  two  atoms : 


nN.. 


1 vulumu. 


3 volumes. 


2 volumes. 


GENERAL  PROPERTIES  OF  GASES. 


77 


In  the  same  way  it  may  be  shown  that  the  phosphorus  molecule  consists 
of  at  least  four  atoms  of  phosphorus  (P4),  etc.,  etc. 

Many  other  facts  corroborate  the  assumption  that  the  molecules  of  the 
elements  consist  of  several  atoms,  for  example,  the  existence  of  the  allo- 
tropic  modifications  of  the  elements  (compare  Ozone),  the  chemical  reac- 
tions (compare  Hydrogen  Peroxide),  and  the  remarkable  action  of  the 
elements  in  the  moment  of  their  liberation. 

We  saw  (p.  53)  that  the  oxygen  separated  from  water  by  chlorine 
acted  much  more  energetically  than  free  oxygen.  Other  elements,  espe- 
cially hydrogen,  behave  similarly  in  the  moment  of  fomnation — in  statu 
nascendi.  As  viewed  by  the  atomic  molecular  theory,  this  may  be  very 


Fig.  38. 


Fig.  37. 


Fig.  39. 


easily  explained.  The  free  elements  (their  molecules)  are  compounds  of 
similar  atoms  whose  chemical  affinity  has  already  been  partly  satisfied. 
In  the  moment  of  their  separation  from  compounds  free  atoms  appear, 
which,  before  they  combine  to  molecules,  must  act  more  energetically. 

The  experiments  illustrating  and  confirming  what  has  been  said  upon 
the  volume  relations  of  the  gases  will  now  be  considered  : 

I.  The  concentrated  acjueous  solution  of  hydrochloric  acid  is  decomposed  by  the  action 
of  the  galvanic  current,  and  the  chlorine  and  hydrogen  collected  ; these  gases  separate  at 
ojjjxjsite  poles.  The  electrolysis  may  be  made  in  an  ordinary  voltameter  (Fig.  37). 
Hofmann’s  apj)aratus  is  better  adapted  to  this  }>urpose  (Fig.  38  *).  Two  glass  cylinders, 

* I'ijf.  38  represents  Uie  volume  relations  the  g-ases  liberated  in  the  electrolysis  of  water; 
compare  p.  92. 


78 


INORGANIC  CHEMISTRY. 


provi(lo<l  at  llie  lop  vvitli  .sto])-c()cks,  are  connected  at  the  lower  end  with  one  another  and 
with  a funnel  tube  ; the  latter  serves  to  lill  the  a])paratus  with  hijiiid  ; and  also,  hy  further 
additions,  to  jness  out  the  gases  collected  in  the  lubes.  'The  platinum  electrodes  are  fused 
into  the  lower  part  of  both  tubes.  I n another  form  of  I lolinaim’s  apparatus  (l‘ig.  yj) 
the  electrodes  are  introduced  by  means  of  caoutchouc  stoi)i)ers.  When  the  separating 
gases  (in  this  case  chlorine)  attack  the  platinum,  carbon  electrodes  are  substituted  for  the 
latter. 

To  electrolyze  hydrogen  chloride,  fill  the  aj)i)aratus  with  hydrochloric  acid,  of  specific 
gravity  1.125,  to  which  has  been  added  fifteen  times  its  volume  of  calcium  chloride  of 
1.36  si)ecilic  gravity  ; close  the  up])er  sto])-cocks  and  connect  the  electrodes  with  the  poles 
of  the  battery,  (bases  separate  in  both  tubes,  and  in  equal  iwlu/nes  ; that  se])aratcd  at  the 
positive  ])ole  is  chlorine  ; the  other  combustible  gas  is  hydrogen.  [See  Liipke,  (Tund- 
zilge  der  wissensch.  Elektrochemie,  llerlin,  1899;  Lolhar  Meyer,  Her.  27  (1894),  850; 
Haber  and  Grimberg,  Z.  f.  anorg.  Ch.  16  11898).] 

d'liis  experiment  shows  that  hydrogen  chloride  decomposes  into  e(|ual  volumes  of 
chlorine  and  hydrogen. 

2.  The  production  of  liydrogen  chloride  by  the  union  of  equal  volumes  of  hydrogen 
and  chlorine  is  shown  in  the  next  experiment. 

Fill  a cylindrical  glass  tube,  provided  with  stop-cocks  at  both  ends  (hig.  40),  with 
equal  volumes  of  chlorine  and  hydrogen.  This  is  most  conveniently  done  by  conducting 
the  gaseous  mixture  obtained  by  the  electrolysis  of  hydrochloric  acid  into  the  dry  tube. 
(The  tube  should  be  filled  and  kept  in  the  dark,  as  the  gases  combine  in  daylight.) 

When  the  tube  is  filled  with  the  mixture,  sunlight  or  magnesium  light 
is  directed  upon  it,  when  chemical  union  ensues.  On  immersing  the 
lower  end  of  the  tube  into  water,  and  opening  the  lower  stop-cock,  the 
water  will  rapidly  fill  the  tube,  as  the  hydrogen  chloride  that  was  pro- 
duced dissolves  ; all  the  hydrogen  and  all  the  chlorine  have  disappeared. 

3.  A modification  of  this  experiment  teaches  us  another  important 
fact  which  has  reference  to  the  ratio  of  the  volume  of  the  hydrogen 
chloride  to  the  volumes  of  its  constituents.  If  the  tube  filled  with 
equal  volumes  of  chlorine  and  of  hydrogen  be  opened  under  mercury, 
after  the  explosion,  no  diminution  in  volume  will  be  detected,  although 
the  mixture  of  chlorine  and  hydrogen  has  been  changed  to  hydrogen 
chloride.  It  follows  from  this  that  a mixture  of  equal  volumes  of 
chlorine  and  hydrogen  affords  the  same  volume  of  hydrogen  chloride, 
or,  as  ordinarily  expressed,  one  volinne  of  chlorine  and  one  volume  of 
hydrogen  yield  tzvo  volutnes  of  hydrogen  chloride. 

The  following  experiment  confirms  this  conclusion  : Into  a bent  tube 
(Fig.  41),  filled  with  mercury,  conduct  dry  hydrogen  chloride,  and  then 
introduce  in  the  bend  of 
the  upper  part  a little  piece 
of  metallic  sodium.  On 
heating  the  latter  with  a 
lamp,  the  hydrogen  chloride 
is  decomi)osed,  the  chlorine 
combines  with  the  sodium 
to  form  sodium  chloride, 
while  hydrogen  is  set  free. 

Upon  measuring  the  resid- 
Fig.  40.  liydrogen  it  will  be  4** 

found  that  its  volume  is 

exactly  the  half  of  the  volume  of  hydrogen  chloride  originally  introduced.  In  the  same 
manner  may  be  shown  the  fact  that  in  two  volumes  of  hydrogen  bromide  and  hydrogen 
iodide  there  is  contained  in  each  one  volume  of  hydrogen.  It  follows  further  from  the 
(hmsities  of  bromine  and  iodine  vaj^ors,  that  the  (juantilies  of  these  elements  in  gas  form 
combinitig  with  one  volume  of  hydrogen  also  occupy  one  volume.  Hence,  one  volume 
of  hydrogen  and  one  volume  of  bromine  7>afor  yield  /zoo  zwlumes  of  hydrogen  bromide, 
and  one  zwlume  of  hydrogen  and  one  volume  of  iodine  vapor  tzvo  volumes  of  hydrogen 
iodide. 


GENERAL  PROPERTIES  OF  GASES. 


79 


All  that  has  been  j^-eviously  stated  may  be  summarized  as  follows: 

1.  A//  bodies  ewe  coinposcd  of  atoms. 

2 . The  atoms  unite  to  foriji  the  7noiccuies  of  the  simpte  and  compound  bodies. 

3.  Equai  votiwies  of  gases  contain,  at  tike  teniperatu7'e  and  under  tike 
pressure,  an  equat  number  of  motecutes. 

4.  The  gas  densities,  therefore,  bear  the  sanie  ratio  to  one  another  as  the 
motecuiar  weights. 

The  gases  argon  and  heiiuin,  recentiy  discovered  in  the  atmosphere,  are 
exceptions.  They  co?isist  of  singie  atoms  at  the  ordinary  temperature. 
This  is  aiso  true  of  the  haiogens  at  etevated  temperatures,  and,  so  far  as  we 
know,  of  metai  vapors  {inercury,  cad7nium,  zi7ic'). 

Heretofore  gas  density  has  been  referred  to  the  hydrogen  atom  (H  = 
i),  while  the  molecular  weight  was  referred  to  the  molecule  H2  = 2,  so 
that  the  density  (the  specific  gravity  of  the  gas)  was  one-half  of  the  molec- 
ular weight.  At  present,  the  atomic  weight  of  oxygen  (16)  being  taken 
as  the  standard  in  determining  the  atomic  numbers  of  the  other  elements, 
it  appears  proper  to  refer  the  density  to  the  molecular  weight  of  oxygen, 
Og  = 32,  whe7i  the  vatues  for  density  and  7noiecuiar  weight  wiil  coincide. 

The  following  table  contains  the  atomic  and  molecular  weights  of  cer- 
tain metalloids  as  well  as  the  molecular  weights  of  their  hydrogen  deriva- 
tives. The  molecular  weights  (gas  densities)  show  how  much  a volume 
of  the  respective  gas  weighs,  if  an  equally  large  volume  of  oxygen,  under 
similar  external  conditions,  be  placed  at  32  (p.  73)  : 


Atomic  Weights. 

Molecular  Weights  or  Gas  Densities. 

H = 1. 01 

U, 

— 

2.02 

ci=  35.45 

Cb 

70.9 

Br  79.96 

Br^ 

= 

159.92 

Br 

79.96 

above 

1000° 

I = 126.85 

I2 

= 

253.70 

I 

= 

126.85 

above 

1500° 

HCl 

= 

36.46 

HBr 

80.97 

HI 

127.86 

0 =r  16 

O2 

32 

II 

H,0 

N, 



18.02 

28.08 

NH3 

= 

17.07 

P = 31.0 

= 

124.00 

PH3 

34.03 

As  one  liter  of  oxygen,  under  normal  conditions  (p.  44),  weighs  1. 4291  grams,  the 
weight  of  one  liter  (L)  of  a gas  of  molecular  weight  M can  be  calculated  from  the  equation 

I,  = = 0.04466  M. 

32 

The  density  D of  the  gas,  referred  to  air  as  unit,  follows  from  the  equation 

1)  = -^- 
28.95 

because  the  density  of  air  is  4291^^  ~ 28.95,  referred  to  Oj  = 32. 


So 


INORGANIC  CHEMISTRY. 


The  numbers  deduced  from  tlicse  equations  vary  somcwliat  from  tliose  obtained  by 
direct  observation,  d'liis  is  due  to  the  fact  tliat  wluit  has  been  said  is  a simple  but  at  the 
same  time  not  an  absolutely  correct  expression  of  the  dej)ortment  of  gases.  W'e  shall 
return  to  this  point  when  the  measurement  of  gases  is  discussed. 

by  determining  the  gas  density  it  is  only  possible  to  fix  the  molecular  weight  of  an 
element  which  exists  in  the  gaseous  condition,  d'he  magnitude  of  the  atomic  weight 
remains  uncertain.  'I'he  chlorine  molecule  weighs  from  its  gas  density  and  the  standard 
selected  70.9  units,  and  consists  of  two  atoms  (Cl.^),  if  we  suppose  that  the  atomic  weight 
= 35-45-  Its  atomic  weight  could,  however,  be  only  the  half  (or  another  submultiple) 
of  35.45  ; then  its  molecule  would  consist  of  four  chlorine  atoms  (Cl^  — 70.90  when  Cl  is 
made  equal  to  17.725),  and  the  formula  of  hydrogen  chloride  would  be  IICl.^.  As  anal- 
ysis and  the  vapor  density  determination  of  many,  especially  organic,  derivatives  show 
that  the  smallest  quantity  of  chlorine  in  a molecule  of  a gasifiable  body  must  always  be 
expressed  by  35.45,  we  can  well  accept  this  number  as  the  atomic  weight.  'I'hat  the 
maximum  values  thus  derived  have  not  been  found  too  high,  but  correspond  to  the  actual 
relative  atomic  weights,  follows  from  the  agreement  of  these  numbers  with  the  atomic 
numbers  obtained  from  the  specific  heat  of  the  elements.  The  complete  certainty  of  their 
correctness  we  reach  by  the  law  of  periodicity,  which  is  deduced  from  these  numbers. 


These  convincing  suppositions  and  conclusions  drawn  from  actual  re- 
lations, form  the  atomic  molecular  doctrine,  which  is  the  foundation  of 
the  chemistry  of  to-day.  As  this  doctrine  comjtletely  explains  the  quan- 
titative phenomena  arising  in  the  action  of  the  chemical  elements  upon 
one  another,  and  as  it  has  been  repeatedly  confirmed  by  phenomena  of  a 
purely  physical  nature,  and  has  thereby  acquired  a high  degree  of  proba- 
bility, it  is  only  proper  and  correct  that  it  be  designated  a theory. 

■ 

OXYGEN  GROUP. 

In  this  group  are  included  the  elements  oxygen,  sulphur,  selenium,  and 
tellurium.  They  are  perfectly  analogous  in  their  chemical  deportment. 
One  atom  of  each  of  these  elements  is  capable  of  uniting  with  two  atoms 
of  hydrogen. 

1.  OXYGEN. 

Atom  : O = 16.  Molecule : O2  = 32. 

Oxygen  (oxygenium)  is  the  most  widely  distributed  element  in  nature. 
It  is  found  free  in  the  air;  in  combination  it  exists  in  water.  It  is  an 
iin])ortant  constituent  of  most  of  the  mineral  and  organic  substances. 

It  was  discovered,  almost  simultaneously,  by  Priestley,  in  England, 
i77<i',  and  Scheele,  in  Sweden,  1775.  Lavoisier,  in  France,  1774-1781, 
first  exjilained  the  imi)ortant  role  attached  to  oxygen  in  processes  of  com- 
bustion, of  res])iration,  and  of  oxidation. 

Pj'eparation.  — Heat  red  mercuric  oxide,  a compound  of  mercury  with 
oxygen,  ima  small  glass  retort;  in  this  way  the  oxide  is  deconq)osed  into 
mercury  and  gaseous  oxygen  : 

llg()=:  Ilg  + O. 


OXYGEN. 


8l 


The  following  method  is  commonly  pursued  in  the  chemical  labora- 
tory : Potassium  chlorate,  a compound  of  potassium,  chlorine  and  oxy- 
gen, is  heated  in  a glass  retort  (Fig.  42)  or  flask,  and  thus  decomposed 
into  solid  potassium  chloride  and  oxygen  : 

KCIO3  = KCl  -f  3O.* 

The  evolution  of  the  gas  proceeds  more  regularly  and  requires  a less 
elevated  temperature  if  the  pulverized  chlorate  be  mixed  with  ferric  oxide 
or  manganese  peroxide.  (A  small  portion  of  the  oxygen  is  ozonized  ; see 
p.  84).  The  liberated  oxygen  is  collected  over  water. 

McLeod  (Jr.  Chem.  Soc.,  55,  192)  explains  the  mechanism  of  the  action  of  man- 
ganese peroxide  on  potassium  perchlorate,  when  heated,  as  follows  : First,  the  peroxide 
acts  on  the  chlorate  producing  potassium  permanganate,  chlorine,  and  oxygen,  2Mn02  + 
2KCIO3  = K2Mn203  -f-  CI2  -f-  O2.  Secondly,  the  permanganate  is  decomposed  by  heat : 
K2Mn208  = K2Mn04  + Mn02  + O2,  and  in  the  third  stage  the  change  is  probably : 
K2Mn04  -|-  CI2  = 2KCI  -|-  MnOg  + 62. 

Very  pure  oxygen  may  also  be  obtained  by  heating  potassium  bichro- 
mate with  concentrated  sulphuric  acid  : 

K2Cr20^  + SHaSO^  = Cr2(SOj3  + 2KHSO4  + 4H2O  + 3O. 

Besides  these  many  other  methods  may  be  employed  for  the  prepara- 
tion of  the  gas:  e.  g.,  the  ignition  of  manganese  and  barium  peroxides; 
the  decomposition  of  sulphuric  acid  at  a high  heat ; the  boiling  of  a solu- 
tion of  bleaching  lime  with  a cobalt  salt,  etc.  These  methods,  applied 
technically,  will  be  considered  more  fully  later. 

A very  convenient  laboratory  method  for  the  preparation  of  oxygen  consists  in  allow- 
ing dilute  hydrochloric  acid  to  act  upon  a mixture  of  barium  peroxide  (2  parts)  and  man- 
ganese peroxide  (l  part).  The  gas  is  evolved  at  the  ordinary  temperatures.  If  the  solid 
ingredients  are  mixed  with  gypsum  (l  part)  and  a little  water  the  mass  can  be  moulded 
into  cubes,  and  the  oxygen  then  be  generated  (like  chlorine)  in  a Kipp  apparatus  (Ber. 
20  (1887),  1585)*  Kassner  obtains  a regular  and  continuous  flow  of  the  gas  by  pouring 
water  upon  r,  pulverized  mixture  of  i part  of  barium  peroxide  and  2j4  parts  of  potassium 
ferricyanide  (red  prussiate  of  potash)  : 

Ba02  -f  K3BaFe2(CN),2  + O3. 

Elkan,  in  Berlin,  furnishes  oxygen  under  100  atmospheres  pressure  in  steel  cylinders  for 
trade  purposes.  The  gas  is  prepared  by  the  method  of  Boussingault  and  Brin,  by  heat- 
ing barium  peroxide  to  about  700°  under  reduced  pressure  : 

Ba02  = BaO  -f  O. 

The  resulting  barium  oxide  is  reconverted  into  peroxide  at  the  same  temperature  but 
higher  pressure  by  the  oxygen  of  the  air.  According  to  Kassner,  calcium  plumbate  may 
be  similarly  applied.  It  breaks  down,  under  conditions  which  will  be  considered  later, 
into  lead  oxide,  calcium  oxide,  and  oxygen  : 

Ca2PbO^  BbO  -f  2CaO  -f  O. 


* The  chemical  equations  used  here  and  ])reviously  are  only  intended  to  represent  the 
manner  of  the  reaction,  and  to  ex])ress  the  accom])anying  relative  (juan titles  by  weight. 
It  should  not  be  forgotten  that  free  atoms  do  not  exist,  but  that  they  always  occur  com- 
bined in  molecules.  Molecularly  written  the  equation  would  be  : 

2KCIO3  = 2KCI  -f  3O2. 


82 


INORGANIC  CIIFMISTRY. 


A method  wliich  will  jwohahly  he  of  value  later  is  that  of  Linde,  hy  which  iiitropjen  is 
first  evaporated  from  litjuid  air,  and  the  portions  passinj^  over  suhsecjuently  are  lound  to 
contain  75  per  cent,  of  oxygen. 

Properties. — Oxygen  is  ca  colorless,  odorless,  tasteless  gas.  One  liter 
of  oxygen  at  0°  C.,  and  760  niin.  jjressiire,  weighs  1.4291  grains  (15.9 
times  more  than  one  liter  of  hydrogen).  Its  sjtecific  volume  ecpials  699.  7 
c.c.  (}).  45).  It  is  only  slightly  soluble  in  water;  100  volumes  of  the 
latter  dissolve  4.1  volumes  of  the  gas  at  0°,  and  2.9  volumes  at  15°.  It 
is  more  readily  dissolved  by  absolute  alcohol  (28  volumes  in  100  volumes). 

The  critical  temiterature  of  oxygen  is — 118°,  and  its  critical  jtressure 
equals  50  atmosi)heres  (]:>.  47).  Idtiuid  oxygen  under  a pressure  of  i 
atmosphere  boils  at  — 184°,  and  under  9 mm.  j)ressure  at  — 225°.  Its 
specific  gravity  at  — 118°  equals  0.65,  at  — 139*^  it  is  0.87,  and  1. 124  at 
— 184°  (the  boiling  point  at  the  ordinary  pressure).  Liquid  oxygen  has 
a bright-blue  color  (Olszewski). 

Oxygen  combines  with  all  the  elements  excepting  fluorine,  helium,  and 


Fig.  42. 

argon.  With  most  of  them  it  unites  directly,  accompanied  by  the  evolu- 
tion of  light  and  heat  (p.  57).  The  burning  of  combustible  bodies  in 
the  air  depends  on  their  union  with  oxygen,  which  is  present  in  the  same 
to  the  amount  of  23  per  cent.  The  phenomena  of  the  respiration  of  ani- 
mals are  also  influenced  by  the  contact  of  the  oxygen  of  the  air — hence 
the  earlier  designations  of  oxygen  as  inflammable  air,  and  vital  air.  In 
j)ure  oxygen  the  phenomena  of  combustion  proceed  more  energetically 
than  in  air.  Ignited  charcoal  or  an  ignited  sliver  inflames  immediately 
in  the  gas,  and  burns  with  a bright  light.  This  test  serves  for  the  recog- 
nition of  pure  oxygen.  Sulphur  and  phosphorus  ignited  in  the  air  burn  in 
it  with  an  intense  light.  Even  iron  burns  in  the  gas.  To  execute  this 
experiment,  take  a steel  watch  spring,  previously  ignited,  attach  a match 
to  the  end,  ignite  the  same,  and  then  introduce  the  spring  into  a vessel 
filled  with  oxygen  gas.  At  once  the  match  inflames  and  ignites  the  iron, 


OXYGEN. 


83 


which  burns  with  an  exceedingly  intense  light  and  emits  sparks.  (To 
protect  the  vessel  from  the  fusing  globules  of  iron  oxide,  cover  the  bottom 
with  a layer  of  sand.)  Iron  will  burn  in  any  flame  if  a current  of  oxygen 
be  conducted  into  the  same. 

Oxygen  combines  with  hydrogen  to  form  water.  The  union  occurs  at 
a red  heat,  by  the  electric  spark  or  by  the  action  of  platinum  sponge  (p. 
46).  Palladium  asbestos  acts  similarly.  Of  all  the  combustible  gases 
only  hydrogen  combines  by  its  influence  with  oxygen  at  the  ordinary 
temperature.  Hydrogen  burns  in  oxygen  with  a flame;  vice  versa,  oxy- 
gen must  also  burn  in  hydrogen;  this  may  be  demonstrated  in  the  same 
manner  as  indicated  under  Hydrogen  Chloride  (p.  57).  A mixture  of 
hydrogen  and  oxygen  detonates  violently ; most  strongly  if  the  propor- 
tions are  i volume  of  oxygen  and  2 volumes  of  hydrogen ; such  a mix- 
ture is  known  as  oxyhydrogen  gas.  The  explosibility  may  be  shown  in 
a harmless  way  by  the  following  experiment:  Fill  a narrow-necked  flask 
of  150-250  c.c.  capacity,  over  water,  yi  with  hydrogen,  and  ^ with 
oxygen;  close  the  opening  with  a cork,  then  wrap  the  flask  in  a towel, 
remove  the  cork  and  bring  a flame  near  the  opening.  A violent  explo- 
sion ensues,  generally  with  complete  breaking  of  the  flask.  The  tempera- 
ture of  ignition  of  oxyhydrogen  gas  is  about  650°  (V.  Meyer  and 
A.  Miinch,  Ber.  26  (1893),  2421)  (see  p.  28). 


Fig.  43. 


The  oxyhydrogen  flame  is  only  faintly  luminous ; it  possesses,  how- 
ever, a very  high  temperature,  answering,  therefore,  for  the  melting  of 
substances  which  fuse  with  great  difficulty,  e.  g.,  platinum.  To  get  a 
continuous  oxyhydrogen  flame,  efflux  tubes  of  peculiar  construction  are 
employed  (Fig.  43)  ; through  the  outer  tube,  W,  hydrogen  is  brought 
from  a gasometer;  oxygen  is  conveyed  through  the  inner,  S,  and  the 
mixture  ignited  at  a.  Such  a flame  impinging  on  a piece  of  burnt  lime 
or  zircon  makes  the  latter  glow  and  emit  an  extremely  bright  light — 
Dru77ii7iond' s Iwie  light,  Zirco7tiu77i  light. 


The  union  of  oxygen  with  other  substances  is  termed  oxidatio7t.  This 
term,  as  well  as  the  name  Oxygenium  (from  and  yewaM)),  or  acid 
producer,  suggested  by  Lavoisier,  arises  from  the  fact  that  acids  are  some- 
times formed  in  oxidation.  This  the  combustion  experiments  prove. 
If  the  vessels,  for  instance,  in  which  carbon,  sulphur,  and  phosphorus 
were  burned,  be  shaken  wu'th  water,  the  latter  will  give  an  acid  taste,  and 
redden  blue  litmus-paper.  It  was  formerly  thought  that  the  formation 


84 


inorc;anic  chemisirv. 


of  acids  is  always  conditioned  by  oxygen.  We  liave,  however,  already 
noticed  that  the  haloid  acids — hydrochloric,  hydrobronhc,  and  hydriodic 
— contain  no  oxygen.  Some  of  the  elements  yield  acids  by  their  union 
with  oxygen,  or  more  correctly  oxides,  which  form  acids  with  water.  Most 
of  these  are  the  metalloids.  Tims  the  following  corresi)onding  acids  are 
derived  from  the  acid-forming  oxides  of  suli)hur  and  phosphorus: 

SO3  + II/)  = II,S(),. 

Sulphur  Sulpliuric 

trioxide.  acid. 

r,/)^  -f  1 1^0  = 2iiro3. 

Phosphorus  Mctaphosptioric 

pentoxide.  acid. 

With  oxygen  the  metals  usually  yield  oxides,  which  form  hydroxides 
(hydrates)  or  bases  with  water  : 

K,0  -f  lIjO  = 2KOII. 

Potassium  Potassium 

oxide.  hydroxide. 

CaO  -f  II2O  = Ca  (011)2. 

Calcium  Calcium 

oxide.  hydroxide. 

The  salts  are  produced  by  the  alternating  action  of  acids  and  bases 
(see  p.  61). 

Thirdly,  there  exist  the  so-called  iiidiffereJit  oxides,  which  yield  neither 
acids  nor  bases,  with  water,  e.  g.  : 

N2O  NO  Pb02. 

Nitrous  Nitric  Lead 

oxide.  oxide.  peroxide. 

Oxidation  is  not  only  induced  by  free  oxygen  or  bodies  rich  in  it,  but 
frequently,  also,  by  the  halogens;  in  the  latter  case  the  halogens  first 
decompose  the  water  with  the  elimination  of  oxygen,  which  then  oxidizes 
further  (compare  p.  53). 

The  opposite  of  oxidation,  the  removal  of  oxygen,  is  called  reduction 
(p.  64).  Hydrogen  {i?i  statu  nascendi'),  and  substances  giving  it  off 
easily  (as  hydriodic  acid),  have  a reducing  action.  Most  of  the  metallic 
oxides  are  reduced  at  a red  heat  by  hydrogen,  with  the  formation  of 
water,  e.  g.  : 

CuO  + H2  = Cu  + II2O. 

Copper  oxide.  Copper. 


OZONE. 

O3. 

Ozone,  discovered  in  1840,  by  Schonbein,  is,  as  will  be  shown  later,  a 
peculiar  modification  of  oxygen,  characterized  by  a remarkable  odor  and 
great  reactivity,  therefore  it  is  called  active  oxygen.  It  is  obtained  from 
oxygen  in  various  ways;  it  is  almost  always  ])roduced  when  oxygen 
is  liberated,  or  when  it  takes  ])art  in  a reaction  ; thus,  in  the  decomposi- 
tion of  peroxides  by  concentrated  sulphuric  acid,  or  when  thevare  heated 
in  a current  of  oxygen  to  their  decomposition-temperature  [Brunk,  Z.  f. 
anorg.  (diem.  10  (i^^95),  222]  ; in  the  electrolysis  of  water  (at  the  positive 


OZONE. 


85 


pole),  in  the  slow  oxidation  of  moist  phosphorus,  in  the  combustion  of 
hydrocarbons,  and  in  the  action  of  the  so  called  silent  discharge  in  an 
atmosphere  of  oxygen  or  air.  In  none  of  these  instances  is  all  of  the 
oxygen  ever  converted  into  ozone  ; only  a small  portion — under  most 
favorable  conditions  5-6  per  cent. — suffers  this  change. 

The  following  methods  serve  for  the  preparation  of  ozone : 

1.  Put  several  pieces  of  stick  phosphorus  into  a spacious  flask,  cover  them  about 
half  with  water,  and  allow  them  to  stand  for  some  hours.  Or  conduct  oxygen  over 
pieces  of  phosphorus  placed  in  a glass  tube  and  moistened  with  water.  Ozone  is  also 
formed  abundantly  when  a potassium  bichromate  solution  is  substituted  for  water. 
[Leeds,  Ann.  Chem.  ig8  (1879),  38.] 

2.  Pass  the  silent  discharge  through  air  or  oxygen.  For  this  purpose  we  may  employ 
a Siemen’s  ozone  tube  (Fig.  44,  Geissler  form),  which  consists  of  two  glass  tubes  a and 
b fitting  into  one  another.  Each  is  coated  on  the  inner  side  with  tin  foil  or  gold  leaf. 
The  one  coating  is  connected  with  the  positive  and  the  other  with  the  negative  pole  of 
an  induction  apparatus.  Oxygen  or  air  circulates  between  the  two  tubes  in  the  direction 
indicated  by  the  arrows. 

3.  Gradually  add  barium  peroxide  in  small  portions  (or  potassium  permanganate)  to 
cold  sulphuric  acid  : 

Ba02  + H2SO,  = BaSO,  + H2O  + O. 

The  escaping  oxygen  is  tolerably  rich  in  ozone,  and  is  collected  over  water. 

Ozone  possesses  a highly  penetrating,  peculiar  odor  (hence  its  name 
from  to  smell),  which  by  prolonged  respiration  produces  bad 

results.  In  a thick  layer,  ozone  shows  a bluish  color.  If  ozonized  air  be 
subjected  to  powerful  pressure  (150  atmospheres)  at  a very  low  tempera- 
ture, or  if  ozonized  oxygen  be  conducted  through  a small  tube  cooled  to 
— 184°  by  boiling  oxygen,  the  ozone  will  condense  to  a liquid  with  an 
indigo-blue  color.  Liquid  ozone,  if  preserved  in  a sealed  tube,  passes 
into  a blue  gas,  that  can  be  again  liquefied  by  chilling  it  with  boiling 
ethylene.  Ozone  is  rather  stable  at  the  ordinary  temperature ; and  at 
more  elevated  temperatures  does  not  revert  to  ordinary  oxygen  as  rapidly 
as  it  was  supposed.  This  was  gathered  from  the  experiments  of  Brunk, 
according  to  whom  ordinary  oxygen  is  ozonized  when  it  is  conducted 
over  manganese  peroxide  heated  to  400°.  According  to  Mailfert’s  experi- 
ments ozone  is  fifteen  times  more  soluble  in  water  than  oxygen.  At  the 
ordinary  temperature  it  takes  up  half  its  volume  of  the  gas.  In  solution 
the  ozone  gradually  reverts  to  oxygen  without  formation  of  hydrogen 
peroxide.  Unlike  ordinary  oxygen,  ozone,  especially  in  a moist  state, 
oxidizes  strongly  at  ordinary  temperatures.  Phosphorus,  sulphur,  and 
arsenic  are  converted  into  phosphoric,  sulphuric,  and  arsenic  acids; 
ammonia  is  changed  to  nitrous  and  nitric  acids  ; silver  and  lead  are  con- 
verted into  the  corresponding  peroxides;  therefore  paper  moistened  with 
a lead  salt  is  colored  brown.  Iodine  is  separated  at  once  from  potassium 
iodide  by  it : 

2KI  -f  II.2O  -f  O3  =:  2K0n  + I2  + O2. 

It  also  oxidizes  all  organic  substances  ; therefore  the  apparatus  used  in 
its  preparation  must  not  be  constructed  of  caoutchouc.  Solutions  of 
dyestuffs,  like  indigo  and  litmus,  are  decolorized.  Its  ability  to  turn  an 
alcoholic  solution  of  guaiacum  tincture  blue  is  very  characteristic  of  ozone. 


86 


INORGANIC  CHEMISTRY. 


For  the  detection  of  ozone  the  ordinary  potassium  iodide  starch-paper  (Schdnhein) 
may  be  used.  This  is  pre])ared  by  immersing  wliite  tissue-paper  in  a starch  solution 
mixed  with  j)otassium  iodide.  Tlie  iodine  which  the  ozone  liberates  from  the  potassium 
iodide  blues  the  starch-paper.  I'he  (piantity  of  ozone  may  be  approximately  determined 
from  the  rapidity  and  the  intensity  of  the  coloration,  'rhallous  hydroxide  is  a more 
reliable  reagent  for  ozone  than  the  ])otassium  iodide  starch-paj)er.  (iuaiacum  tincture  and 
paper  saturated  with  a lead  acetate  .solution  may  also  l)e  used  to  detect  ozone  ; the  hrst 
ac(|uire.s  a blue  color,  the  second  becomes  brown.  Other  substances  al.so  blue  potassium 
iodide  starch-paper  and  guaiacum,  e.  g.,  chlorine,  bromine,  nitrogen  dioxide,  etc.,  etc.  d'o 
distinguish  ozone  from  the.se,  proceed  as  follows  (Ilouzeau)  : 'I'ake  two  strips  of  violet 
litmus-i)aper,  one  of  which  is  saturated  with  jjotassium  iodide,  and  exj)ose  it  to  the  action 
of  the  gas  ; when  ozone  is  present  potassium  hydroxide  will  be  formed  from  the  iodide,  atid 
color  the  violet  litmus  blue.  'Fhe  second  paper  serves  to  show  the  ab.sence  of  ammonia. 

'rhe  preceding  reactions  of  ozone  are  all  produced  by  hydrogen  peroxide,  although 
less  ra])idly.  The  only  test  answering  for  the  distitiction  of  very  slight  (piantities  of 
ozone  from  hydrogen  peroxide  is  the  blackening  of  a bright  strip  of  silver  by  ozone. 


Fig.  44. 


Ozone  is  a variety  of  oxygen its  molecules  consist  of  three  atoms: 

3O2  yield  2O3. 

3 vols.  oxygen.  2 vols.  ozone. 

This  is  })roved  by  the  following  experiments  : In  ozonizing  oxygen  its  volume  dimin- 
ishes ; upon  heating  (when  ozone  is  again  changed  to  oxygen),  the  original  volume  is  repro- 
duced ; when  ozonized  oxygen  is  brought  in  contact  with  oil  of  turpentine  or  oil  of  cinna- 
mon, all  the  ozone  is  absorbed  and  the  volume  of  the  gas  is  diminished.  Comj)ari!ig  this 
diminution,  corresponding  (o  the  ozone  volume,  with  the  expansion  which  an  equal  vol- 
ume of  ozonized  oxygen  suffers  after  the  application  of  heat,  we  will  find  that  the  first 
is  twice  as  large  as  the  latter  ; this  indicates  that  I volume  of  ozone  yields  i ^ volumes  of 
oxygen  (.Sorc't).  h'rom  this  it  follows  that  the  s|)ecific  gravity  of  ozone  must  be  I times 
grcaU*r  than  that  of  ordinary  oxygen,  and  that  if  the  molecule  of  oxygen  consists  of  two 
atoms,  the  molecule  of  ozone  must  contain  three  atoms.  'I  his  conclusion  is  confirmed  by 
the  specific  gravity  of  ozone  derived  experimentally  from  the  velocity  of  diffusion,  and  has 


OZONE.  87 

been  found  to  be  approximately  48  (O2  = 32),  corresponding  to  the  molecular  formula 
O3  (Graham). 

A diminution  in  the  volume  of  the  gas  does  not  occur  in  the  action  of 
ozone  upon  oxidizable  bodies  like  potassium  iodide  and  mercury,  although 
all  the  ozone  disappears.  It  would  appear  from  this  that,  in  oxidizing, 
ozone  only  acts  with  one  atom  of  oxygen,  while  the  other  two  atoms  form 
free  oxygen,  which  occupies  the  same  volume  as  the  ozone : 

O3  -|-  2K.I  = O2  ”b  K2O  -j-  l2« 

I vol.  I vol. 

As  a consequence  of  this  behavior  ozone  is  also  called  oxidized  oxy- 
gen; i.  e.,  free  oxygen  (O2),  which  has  combined  with  an  additional 
oxygen  atom. 

Thermo -cheviical  Deportment. — Compared  with  ordinary  oxygen,  ozone  is  an  endo- 
thermic compound.  Heat  is  absorbed  in  its  formation  from  oxygen  : 

(02,0)  -f  36.2  Cal.  = O3. 

This  explains  why  ozone  is  produced  with  so  much  difficulty,  and  why  the  addition  of 
considerable  energy  is  necessary.  This  may  be  applied  directly  in  the  form  of  heat  or 
electricity  (electric  sparks,  silent  electric  discharge),  or  it  may  be  withdrawn  from  the 
heat  of  formation  of  other  exothermic  compounds  which  are  produced  at  the  same  time, 
e.  g.,  the  formation  of  ozone  by  the  oxidation  of  phosphorus  to  phosphorous  acid. 

Being  an  endothermic  derivative,  we  readily  perceive  why  ozone  is  so  unstable,  and 
why  it  changes  so  readily  to  ordinary  oxygen.  When  this  occurs  the  oxygen  acts  as  if 
in  the  moment  of  formation  (O3  = O2  + O ; see  p.  53),  and  this  explains  why  ozone 
acts  more  powerfully  than  ordinary  oxygen.  And  to  this  must  be  added  that  in  oxida- 
tions performed  by  ozone  there  are  36.2  Cal.  more  set  free  than  in  oxidations  with  ordi- 
nary oxygen. 

We  observe,  therefore,  that  the  elementary  substance  oxygen  occurs  in 
free  condition  in  two  different  forms — allotropic  modifications — ordinary 
oxygen  (Og)  and  ozone  (O3).  We  shall  learn  later  that  very  frequently 
substances  of  the  same  elementary  composition  possess  different  physical 
and  chemical  properties;  such  bodies  are  called  isomerides  and  the  phe- 
nomenon isomerism  (lVo<r,  like ; p.ipoq,  part).  The  isomerism  of  the  ele- 
ments was  called  allotropy  {aXkorpimoc;,  differently  formed)  by  Berzelius. 
It  is  accounted  for  (as  in  the  case  of  oxygen  and  sulphur)  by  the  differ- 
ent number  of  atoms  in  the  molecule.  Alterations  in  the  molecular 
energy  relations  are  obviously  of  great  importance.  The  allotropy  of 
oxygen  confirms  the  conclusion  drawn  from  the  gas  densities,  that  the 
molecules  of  the  elements  are  composed  of  atoms. 


We  have  already  seen  that  ozone  is  absorbed  not  only  by  turpentine 
and  cinnamon  oil,  but  also  by  other  ethereal  oils.  These  bodies  are, 
however,  very  slowly  oxidized;  the  oxygen  is  contained  in  them  in  a 
peculiar  condition.  In  this  form  it  acts  upon  some  bodies  like  free 
ozone;  in  other  instances,  the  oxidizing  action  is  rendered  possible  only 
by  peculiar  substances  which  carry  the  oxygen.  Spongy  platinum,  ferrous 


88 


INORGANIC  CHEMISTRY. 


sulphate,  and  the  blood-corpuscles  are  examples  of  this  class.  Thus,  old 
turpentine  oil,  containing  absorbed  ozone,  only  acts  on  |Kiper  saturated 
with  starch-paste  and  potassium  iodide,  when  a few  drops  of  a ferrous 
sulphate  solution  have  been  added  to  it. 

Since  ozone  is  formed  when  electricity  acts  ui)on  air,  and,  indeed, 
probably,  in  all  oxidation  and  combustion  processes, — and,  further,  j^otas- 
sium  iodidestarch-paper  is  blued  when  exposed  to  the  air, — it  was  believed 
that  ozone  was  a constant  constituent  of  atmosj^heric  air  (i-io  milli- 
grams in  loo  liters  of  air)  ; according  to  recent  investigations  it  is,  how- 
ever, probable  that  the  imagined  ozone  reactions  are  frecpiently  produced 
by  hydrogen  peroxide,  which  is  very  similar  in  reaction  t(j  ozone  (p.  103), 
and  is  almost  constantly  in  the  air  (Schdne).  [Her.  13  (1880),  1503; 
Engler  and  his  co-workers,  ]k*r.  30  (1897),  1669;  31  (1898),  3046.] 

Antozone,  which  was  regarded  as  a third  peculiar  modification  of 
oxygen,  has  been  proved  in  some  cases  to  be  hydrogen  peroxide  and  in 
others  to  be  oxygen  in  statu  nascendi. 

Oxygen  is  taken  up  at  the  ordinary  temperature  by  comparatively  few  bodies.  The 
alkali  metals,  finely  divided  metals,  compounds  of  sulphurous  acid,  phosphorus,  certain 
organic  derivatives,  etc.,  are  .such  “ auto-oxidizable  ” bodies.  In  their  spontaneous  oxi- 
dation the  very  remarkable  phenomenon  has  been  noticed  that  other  substances,  present 
at  the  same  time  and  in  themselves  not  oxidizable,  are  also  oxidized.  It  was  therefore 
said  that  oxygen  was  rendered  active  by  the  “ auto-oxidizable  ” .substances.  More  recent 
experiments  have  proved  that  the  oxidation  was  occasioned  by  intermediate  products  rich 
in  oxygen  (peroxides)  which  readily  parted  with  their  oxygen  and  acted  as  oxygen 
carriers.  [See  Engler  and  Wild,  etc.] 


COMPOUNDS  OF  OXYGEN  WITH  HYDROGEN. 

1.  WATER. 

Molecule  : H2O  = 18.02. 

Water,  the  product  of  the  union  of  hydrogen  and  oxygen,  is  produced 
in  many  chemical  processes,  e.  g.,  in  the  formation  of  salts  from  bases  and 
acids  (p.  61).  Cavendish  was  the  first  (1781)  to  produce  water  syntheti- 
cally by  the  combustion  of  hydrogen  in  oxygen.  Lavoisier  a little  later 
showed  by  analysis  that  water  contained  hydrogen  and  oxygen  (p.  40). 
The  first  determination,  although  far  from  accurate,  of  the  quantitative 
composition  of  water  also  originated  with  Lavoisier  (1783).  The  weight 
ratio  of  hydrogen  and  oxygen  in  it  was  first  determined  more  correctly 
by  Berzelius  and  Dulong  (1820).  Gay-Lussac  conjointly  with  Humboldt 
(1805)  had  shown  that  water  was  produced  by  the  union  of  two  volumes 
of  hydrogen  with  one  volume  of  oxygen  (p.  74). 

Physical  Properties. — It  is  obtained  chemically  pure  by  the  distillation 
of  naturally  occurring  water,  which  always  contains  other  matter  dissolved 
in  it.  It  apj)cars  in  all  three  states  of  aggregation  : in  the  liquid,  gaseous 
(steam),  anil  s(;lid  (ice,  snow).  In  a thin  layer  water  is  colorless,  while 


WATER. 


89 


it  is  blue  in  layers  of  greater  thickness.  When  water  is  cooled  it  con- 
tracts and  attains  its  greatest  density  at  +3.98°  C.  The  weight  of  a 
cubic  centimeter  at  -1-4°  is  taken  as  the  unit  of  weight  (=  i gram).  By 
further  cooling  the  water  expands — the  opposite  of  most  other  bodies  ; its 
volume  becomes  greater,  while  the  specific  gravity  decreases. 

The  following  table  gives  the  volume  and  specific  gravity  of  water  for 
different  temperatures  referred  to  water  at  4°  (see  the  investigations  of 
Thiesen,  Scheel  and  Diesselhorn,  Zeit.  f.  Instrumentenkunde,  17  (1897), 

332): 


Temperature. 

Specific  Volume. 

Specific  Gravity. 

0° 

1. 0001 324 

0.9998676 

2° 

1.0000320 

0.9999680 

4° 

1. 0000000 

1. 0000000 

6° 

1.0000320 

0.9999680 

8° 

I.0001241 

0.9998759 

10° 

1.0002730 

0.9997271 

12° 

1.0004756 

0.9995246 

14° 

1.0007292 

0.9992713 

16° 

I.0010314 

0.9989697 

18° 

1. 00 1 2000 

0.9986220 

20° 

1.0017728 

0.9982303 

22° 

1.0022083 

0.9977966 

26° 

1.0032006 

0.9968097 

36° 

1.0063297 

0.9937101 

38° 

1.0070584 

0.992991 1 

50° 

1. 01 200 

0.98813 

100° 

1.04327 

0.95863 

By  cooling,  water  solidifies  to  ice.  The  melting  point  of  ice  is  taken 
as  the  zero  of  the  centigrade  and  Reaumur’s  thermometric  scales.  The 
solidification-temperature  of  water  cannot  be  used  for  this  purpose, 
because  water  at  rest  can  be  chilled  considerably  below  0°  without  freez- 
ing. The  freezing  point  of  ice,  however,  like  all  other  solid  bodies,  is 
constant  at  a definite  pressure.  As  the  latter  rises  the  melting  point  of  ice 
falls.  At  a pressure  of  1000  atmospheres  ice  melts  at — 7°  (Lord  Kelvin). 
Most  other  substances,  unlike  ice,  melt  with  an  increase  in  volume  ; in 
their  case  increase  in  pressure  occasions  a rise  in  the  melting  point 
(Bunsen). 

In  the  conversion  of  water  into  ice,  a considerable  expansion  occurs : 
TOO  volumes  of  water  at  0°  yield  109  volumes  of  ice  at  0°  ; the  specific 
gravity  of  the  latter  is,  therefore,  0.9173,  referred  to  water  at  4°.  This 
is  why  ice  floats  in  water.  Ice  crystallizes  in  rhombohedral  forms  of  the 
hexagonal  system  as  may  be  distinctly  observed  in  snowflakes. 

Different  bodies  require  different  quantities  of  heat  to  bring  them  to 
the  same  temperature,  'fhe  heat  capacity  of  water  is  greater  than  that 
of  all  other  liquid  or  solid  bodies.  It  is  customary  to  take  the  quantity 
of  heat  necessary  to  raise  one  part  by  weight  of  water  from  15°  to 
16°  C.,  as  the  unit  of  heat,  or  calorie  (see  p.  66).  In  the  jiassage  of  a 
8 


90 


INORGANIC  CHEMISTRY. 


licjuid  to  the  solid  state  lieat  is  always  set  free,  while,  on  the  other  hand, 
in  the  fusion  of  the  solid  heat  is  absorbed.  'I’he  latent  heat  of  fusion  of 
I kilogram  of  ice  ecjuals  79  calories;  this  means  that,  for  the  melting  of 
the  ice,  a (juantity  of  heat  is  reejuired  which  is  cajiable  of  raising  79  kilo- 
grams of  water  from  15°  to  16°  (j),  92). 

Water  boils  upon  the  a})plication  of  heat,  and  is  converted  into  steam. 
The  boiling  temperature,  like  that  of  all  other  licpiids,  depends  on  the 
pressure  ; it  is  also  influenced  by  the  substances  dissolved  in  it,  although 
the  temperature  of  the  vapors  is  constant  (at  a given  pressure),  d'he 
tem})erature  of  the  steam  escaping  from  water  at  the  ordinary  pres- 
sure of  760  mm.  (45°  latitude;  sea-level)  is  100°  (=  80°  Reaumur).  At 
680  mm.  barometric  pressure  water  boils  at  96.9°,  and  at  800  mm.  at 
101.4°. 

One  volume  of  water  at  100°  C.  yields  1696  volumes  of  vapor  of  the 
same  temperature.  The  specific  gravity  of  steam  equals  18.02  (03=  32) 
or  = 0.622  (air  = i);  see  p.  79.  One  liter  of  aqueous  vapor 

weighs  0.590  gram  at  100°  under  normal  pressure. 

The  critical  temperature  of  water  (or  its  absolute  boiling  temperature, 
p.  48)  is  +37o°j  its  critical  volume  2.33,  and  its  critical  pressure  195.5 
atmospheres,  i.  e.,  at  370°  the  tension  of  its  vapor  equals  195.5  atmos- 
pheres, and  above  this  temperature  it  can  no  longer  exist  as  a liquid,  but 
only  as  a gas. 

The  vaporization  of  water,  and  of  other  liquids,  occurs  not  only  at  the  boiling  point, 
but  also  at  lower  temperatures.  The  tension  of  the  vapor  is  measured  by  the  height  of 
the  column  of  mercury  which  holds  it  in  equilibrio. 

The  following  table  gives  the  tension  of  aqueous  vapor  for  various  temperatures  (ex- 
pressed in  mercury  levels  at  0°,  45°  latitude,  and  sea-level)  : 


IMPERATURE. 

Tension. 

Temperature. 

Tension. 

—20°  C. 

0.93  mm. 

20°  C. 

17.4  mm. 

— 10°  c. 

2. 15  mm. 

40°  C. 

54.9  mm. 

0°  C. 

4. 6 mm. 

60°  C. 

148.9  mm. 

-f  10°  c. 

9. 1 mm. 

80°  C. 

355.4  mm. 

+15°  C. 

1 2. 7 mm. 

100°  C. 

760.0  mm. 

Moist  gases,  therefore,  occupy  a larger  volume  than  those  which  are  dry.  (See  Air : 
measurement  of  gases). 

The  latent  heat  of  the  evaporation  of  a kilogram  of  water  equals  536.4 
heat  units  at  100°  C. ; i.  e.,  for  the  conversion  of  one  kilogram  of  water 
at  100°  C.  into  vapor  of  the  same  temjierature  a quantity  of  heat  will  be 
absorbed  cajtable  of  raising  536.4  parts  of  water  from  15°  to  16°. 

Jn  consequence  of  the  evaporation  of  water,  the  gases  separating 
from  an  atpieous  solution  are  always  moist.  To  dry  the  same,  conduct 
them  over  substances  which  will  take  up  the  moisture,  e.  s^.,  calcium 
chloride,  stick  ])otash,  sulplmric  acid,  ])hosphoric  anhydride  (compare 
j).  42).  Many  solids  al)stract  moisture  from  the  air  without  chemically 
uniting  with  it;  to  dry  these  let  them  stand  in  an  enclosed  space 
over  suli>huric  acid,  calcium  chloride,  or  phosphorus  pentoxide  (desic- 
cators). 


WATER. 


9 


Natural  Wafers. — As  water  dissolves  many  solid,  liquid,  and  gaseous 
compounds,  all  naturally  occurring  waters  contain  foreign  admixtures. 
The  purest  natural  water  is  rain-  or  snow-water;  it  contains  about  3 per 
cent,  by  volume  of  gases  (oxygen,  nitrogen,  argon,  and  carbon  dioxide), 
and  traces  of  solids  (the  ammonium  salts  of  nitrous  and  nitric  acids). 
If  water  that  has  been  standing  exposed  to  the  air  be  heated,  the  dis- 
solved gases  escape  in  bubbles. 

River  and  spring  waters  contain  solid  constituents  in  widely  varying 
amounts.  Water  having  much  lime  and  gypsum  present  in  it  is  ordinarily 
known  as  hard,  in  distinction  from  soft  water,  which  contains  less  lime 
(see  Calcium  Carbonate).  On  boiling  lime-waters,  most  of  the  impurity 
deposits  out.  Spring  water  generally  contains  in  addition  larger  quan- 
tities of  carbon  dioxide,  which  impart  a refreshing  and  enlivening  taste 
to  it.  Spring  waters  holding  considerable  quantities  of  solid  constitu- 
ents, or  exhibiting  special  healing  properties,  are  called  mineral  waters. 
These  are  distinguished  as  saline  waters  (containing  sodium  chloride), 
magnesian  waters,  sulphur  waters  (hydrogen  sulphide),  acidulated  waters 
(saturated  with  carbon  dioxide),  chalybeate  waters  (containing  iron), 
and  others. 

Sea-water  contains  3.5  per  cent,  of  salts,  chiefly  sodium  chloride  (2.7 
per  cent.). 

To  purify  the  natural  waters  they  are  filtered  (for  the  removal  of 
mechanical  admixtures)  and,  for  chemical  and  pharmaceutical  purposes, 
distilled  {distilled  water)  in  apparatus  of  varying  form. 

Chemical  Properties  of  Water. — Water  is  a neutral  substance,  i.  e.,  it 
possesses  neither  acid  nor  basic  properties.  As  we  have  already  observed 
(p.  61),  it  forms  bases  with  basic  oxides  and  acids  with  acid- forming 
oxides. 

Despite  the  fact  that  the  affinity  of  hydrogen  for  oxygen  is  may  great, 
water  may  be  decomposed  by  many  substances.  At  ordinary  tempera- 
tures, metals  like  potassium,  sodium  and  calcium  decompose  it,  with 
liberation  of  hydrogen  : 

2H2O  -f  Kj  = 2KOH  -f  H2. 

Other  metals  do  not  decompose  it,  except  at  elevated  temperatures. 
Steam  conducted  over  ignited  iron  gives  its  oxygen  to  the  latter,  form- 
ing ferroso-ferric  oxide,  while  hydrogen  is  set  free  (pp.  40,  94)  : 

3Fe  + 4H2O  = FcgCh  -f  4H2. 

This  is  a reversible  reaction.  Many  of  these  oxides  (even  sodium  oxide,  according 
to  Beketoff ) are  again  reduced  by  hydrogen  at  more  elevated  temperatures  : 

1^^304  + 411.!  = Fe3  + 4H2O. 

(See  pp.  52,  68,  95.) 

Chlorine  decomposes  water  in  the  sunlight;  the  decomposition  is  more 
rapid  when  the  vapors  are  conducted  through  heated  tubes  (j).  52)  : 

2IT./)  f 2CI2  = 4TTCI  + O2. 

Many  chemical  reactions  are  only  comj)leted  in  the  presence  of  moisture.  Thus,  the 
metals  are  only  oxidi/.cd  at  the  ordinary  temperature  when  both  oxygen  and  moisture  are 


92 


INOKHANIC  CHKMISTRY. 


l)resent.  Iron  docs  not  rust  in  jicrfeclly  dry  air.  Absolutely  dry  oxyjfcn  docs  not  act 
upon  tlie  sul)stanccs  which  it  ordinarily  attacks  with  great  energy.  I’liosplionis,  carbon 
and  carbon  monoxide  do  not  burn,  or  at  least  but  feebly,  in  perfccily  dry  oxygen.  I try 
liydrogen  chloride  does  not  turn  blue  litmus  red,  etc.  [Dixon,  l>er.  19  (1886),  Kef.  157  ; 
Baker,  ibid.  18  (1885),  Ref.  426;  Lothar  Meyer,  I5er.  19  (1886),  1099;  K.  Otto,  l>er. 
26  (1893),  II,  2050;  Hughes,  ibid,  iv,  863;  Veley,  Her.  29  (1896),  i,  577;  also  Out- 
mann,  Ann.  Chem.  299  (1898),  267  ; see  also  p.  loi.] 

Electrolysis  of  Water. — The  electric  current,  acting  upon  water  acidu- 
lated with  sulpliuric  acid,  decomposes  it  ajiparently  directly  into  its  ele- 
ments. Hydrogen  collects  at  the  negative  pole — the  kathode,  while 
oxygen  ajtpears  at  the  positive  jtole — the  anode. The  volume  of  the 
hydrogen  is  nearly  twice  that  of  the  oxygen  (pp.  76,  98). 

The  electrolytic  decomposition  of  water  is  more  complex  than  is  ordinarily  supposed, 
as  perfectly  pure  water  is  not  ca{)able  of  conducting  the  current,  and  is  conseciuently  not 
decomposed  by  it.  It  is  rather  the  added  sulphuric  acid  which  suffers  the  decomposition 
and  by  means  of  the  water  is  always  reformed  (compare  the  Fdectrolysis  of  Salts). 
Hydrogen  and  oxygen  are  merely  the  end-products  of  this  change.  In  addition  to 
oxygen,  about  one  per  cent,  of  ozone  is  produced  ; further,  sulphur  hejjtoxide  (persid- 
phuric  acid)  and  liydrogen  jieroxide  are  formed  at  the  anode.  Some  hydrogen  peroxide 
is  produced  at  the  negative  pole  (the  kathode)  as  the  result  of  the  union  of  nascent 
hydrogen  with  the  dissolved  oxygen  (p.  loi). 

Therfno-chemical  Deportme7it. — Water  is  formed  from  its  elements  with 
the  liberation  of  much  heat.  57.2  Cal.  are  disengaged  in  the  union  of 
2 grams  of  hydrogen  with  16  grams  of  oxygen  to  produce  aqueous  vajior 
of  100°  (H.^,  O — vapor).  In  the  condensation  of  the  steam  to  water  of 
100°  9.63  Cal.  are  liberated  (—  18  X o-SS^);  this  is  the  latent  heat  of 
evaporation  (see  p.  89).  And  again,  in  the  cooling  of  the  water  through 
every  1°  C.,  Cal.  more  escape;  consequently  in  cooling  from  100° 
to  16°  there  would  be  a liberation  of  84.  ^ = 1.5  Cal.  (see  p 66). 

So  that  in  the  production  of  the  molecular  weight  (18  grams)  of  water  of 
16°  temperature  from  its  elements  there  is  a total  disengagement  of  68.3 
large  calories : 

(Hj,  O — vapor)  = 57.2  Cal.  (Hg,  O — liquid)  = 68.3  Cal. 

The  decomposition  of  water  at  the  ordinary  temperature  by  a metal  only  takes  place 
if  the  heat  of  formation  of  the  oxide  is  greater  than  that  of  the  water.  Sixteen  grams  of 
oxygen  in  their  union  with  hydrogen  liberate  68.3  Cal.,  with  sodium  100  Cal.,  and  with 
cop])er  but  38  Cal.  The  decomposition  of  water  by  sodium  is  an  exothermic  reaction, 
while  that  by  copper  is  endothermic  (see  p.  94). 


* Faraday  introduced  the  following  terms,  which  have  been  universally  adopted  : The 
metal  wires  or  j)lates  by  which  the  current  enters  and  pas.ses  from  electrolytes  are  called 
the  electrodes  {bbbr,  way)  ; the  electrode  by  which  the  current  enters  is  the  and  the 

other  through  which  the  ])ositive  electricity  has  its  exit  and  by  which  the  negative  elec- 
tricity enters  is  the  kathode.  That  part  of  the  electrolyte  jjassing  to  the  anode  where  it  is 
se])aratcd  or  deposited  is  the  anion  ( that  which  migrates  upward — opposite  to  the  current 
of  positive  electricity)  while  the  ])ortion  going  to  tlie  kathode  and  separating  there  is  the 
kation  (migrnling  downwards — with  the  current  of  positive  electricity).  Both  are  called 
ions  (/bora,  wandering,  migrating).  As  a rule  the  kations  in  a compound  are  replaceable 
by  hydrogen  ; tlx.*  anions  are  simple  or  compound  halogenides  (Cl,  SO^)  (compare  the 
F.Iectrolysis  of  Salts  and  b'.lect  roly  tic  I )is.sociation). 


WATER. 


93 


Dissociation  of  Water — Water,  like  otlier  clieiiiical  compounds,  is 
broken  down  into  its  elements  by  heat.  This  was  first  observed  upon 
pouring  molten  platinum  into  water,  when  bubbles  of  oxyhydrogen  gas 
appeared  (Grove).  This  decomposition  of  water  was  first  ascribed  to  a 
catalytic  action  of  the  platinum.  Sainte-Claire  Deville  was  the  first  to 
carefully  investigate  and  explain  the  decomjiosition  phenomena  induced 
by  heat,  thus  disclosing  one  of  the  most  important  chajiters  of  theoretical 
chemistry.  He  proved  that  a decomi)osition  (dissociation)  like  the 
preceding  did  not  take  place  suddenly,  l)ut  gradually;  that  it  advanced 
regularly  with  increasing  temperature,  and  was  limited  by  an  opposing 
combination-tendency  on  the  part  of  the  components.  The  temperature 
at  which  the  decomposition  is  half  finished  is  usually  designated  as  the 
temperature  of  decomposition. 

Deville  illustrated  the  decomposition  of  water  by  the  following  experi- 
ment: Pass  aqueous  vapor  through  a porous  clay  tube,  a,  cemented  into 
a wider  non-permeable  porcelain  tube  heated  to  a white  heat  in  an  oven 


(Fig.  45).*  The  water  suffers  partial  decomposition,  the  lighter  hydrogen, 
which  passes  through  the  inner  tube  into  the  porcelain  tube  more  rapidly 
than  the  oxygen,  escapes  through  the  gas  tube  b.  The  oxygen  escapes 
mainly  through  the  inner  tube  at  a.  A part  of  the  same  diffuses  simul- 
taneously with  the  hydrogen  and  reunites  with  the  latter.  To  avoid  this, 
conduct  a stream  of  carbon  dioxide  through  the  wider  porcelain  tube;  this 
will  carry  out  the  hydrogen  with  it.  The  carbon  dioxide  will  be  absorbed 
by  the  alkali  solution  in  the  collecting  vessel,  and  oxyhydrogen  gas  be 
found  in  the  cylinder.  The  decomposition  of  the  water  commences  at 
about  1000°,  and  is  half  finished  at  about  2500°.  The  quantity  of  gas 
liberated  in  equal  periods  rises  successively  with  the  temperature. 

Many  other  compounds,  as  carbon  dioxide,  hydrogen  chloride,  iodine 
(P-  55))  ammonium  chloride,  ])hosphorus  pentachloride,  etc.,  are  simi- 


* A tube  of  platinum  may  be  well  substitutecl  for  the  porous  clay  tube  ; at  a red  heat 
it  permits  the  passage  of  hydrogen,  but  not  that  of  oxygen  (p.  45), 


94 


INORGANIC  CHKMIS'IRY. 


larly  dissociaLcd  by  heal.  These  are  all  exothermic  compounds,  absorb- 
ing energy  in  their  decomposition,  and  are  therefore  decomposed  but 
gradually,  dej)ending  upon  the  amount  of  energy  imparted  to  them.  In 
these  instances  heat  ojiposes  the  affinity  of  the  various  components,  so 
that  if  the  temperature  be  lowered  there  will  occur  a jiartial  reunion  of 
the  same.  The  siilitting  up  of  the  endothermic  comjioiinds  is  entirely 
different,  e.  g.,  that  of  ])otassium  chlorate,  KC1(),„  into  potassium  chloride 
and  oxygen,  of  ammonium  nitrite,  NH^NO^,  into  water  and  nitrogen,  of 
nitrogen  chloride  into  chlorine  and  nitrogen,  etc.  Heat  is  set  free  in 
the  decomposition  of  these  compounds.  Any  added  or  external  heat 
only  incites  or  brings  on  the  decomjiosition  and  overcomes  the  chemical 
affinity.  Under  some  conditions  there  are  accompanying  explosions; 
there  is  no  reunion  of  the  components  on  lowering  the  temperature. 

The  ex])lanation  of  the  dissociation  plienoinena  is  found  in  tlie  kinetic  tlieory  of  gases 
{Kivt/aii',  motion).  According  to  it  the  heat-motion  of  the  molecules  of  a gas  is  motiem 
in  direct  lines,  jjrogressive  and  with  uniform  velocity.  As  the  temperature  rises  the 
velocity  of  this  motion  increases.  Its  energy  for  the  molecules  of  different  gases  at  any 
given  temperature  is  the  same.  Heavy  molecules  move  corres[)ondingly  slower  than  the 
light  molecules.  The  atoms,  too,  forming  the  molecule,  have  their  own  peculiar  motions, 
which  become  more  energetic  with  the  rise  in  temperature.  By  the  peculiar  motions  of 
the  atoms — in  which  the  center  of  gravity  of  the  molecule  is  not  concerned — the  internal 
arrangement  of  the  molecule  is  affected  ; it  akso  opposes  the  action  of  affinity.  Therefore, 
just  as  soon  as,  by  augmented  temperature,  the  centrifugally  active  energy  of  atomic 
motion  equals  or  exceeds  the  affinity,  the  breaking  down  of  the  molecule  takes  place. 
Further,  as  a consequence  of  irregular  collision,  the  molecules  do  not  all  possess  the 
.same  velocity  at  a given  temperature  ; some  move  more  rapidly,  others  .slower,  than  the 
main  portion  ; the  former  are  warmer  than  the  latter.  Only  the  sum  of  the  existing  forces 
of  ail  the  molecules  is  a constant  quantity  at  every  temperature. 

The  molecules  first  decomposed  are  those  moving  more  rapidly  and  having  a temperature 
beyond  the  average  temperature.  Their  number  increases  with  the  temperature.  Hence, 
it  follows  that  dissociation  also  increases  with  rise  of  temperature.  Compare  in  this 
connection  W.  Nernst  and  A.  Schonflies,  Einfiihrung  in  die  math.  Behandlung  der 
Naturwissenschaften,  2 Aufl.  (1898). 

The  dissociation  of  solids,  which  when  heated  develop  gaseous  ingre- 
dients, is  very  instructive  and  remarkable — for  example,  the  decomposi- 
tion of  calcium  carbonate,  CaCOg,  into  calcium  oxide  and  carbon 
dioxide,  of  sodium  and  potassium  hydrides  (K^H2)  into  their  elements, 
etc.  These  indicate  that  the  dissociation  is  not  only  dependent  upon  the 
temperature  but  also  upon  external  pressure,  and  that  for  every  tempera- 
ture there  is  a corresponding  definite  tension  of  dissociation — a pressure 
below  which  the  decomposition  will  not  occur.  For  further  particulars 
on  this  point  see  Potassium  Hydride  and  Calcium  Carbonate. 

Dissociation,  i.  e.,  the  partial  decomposition  increasing  with  the  tem- 
])erature,  ex])lains  many  chemical  processes  and  jffienomena.  Thus  it 
accounts  for  the  abnormal  vapor  densities,  which  apparently  contradicted 
the  law  of  tlie  efjnal  number  of  molecules  being  ])resent  in  equal  volumes 
of  gases  (j).  79)  ; all  variations  from  it  are  always  due  to  the  breaking  down 
of  more  complex  molecules  (see  Sulphur,  p.  106).  The  observed  vapor 
density  affords  a cine  to  the  magnitude  of  the  dissociation.  The  mass 
action  in  reversible  (inverse)  chemical  reactions  is  afforded  a simple  ex- 
})lanation  by  dissociation.  We  have  already  said  that  iron  raised  to  a red 


THE  QUANTITATIVE  COMPOSITION  OF  WATER.  95 

lieat  decomposed  water  with  the  separation  of  hydrogen  and  the  produc- 
tion of  ferrous-ferric  oxide : 

3Fe  + 4H2O  = Fe30^  + 4H,. 

On  conducting  hydrogen  over  ignited  iron  oxides  the  opposite  process 
occurs;  the  oxygen  compound  of  the  iron  is  reduced  and  water  and  iron 
are  formed : 

Fep^  + 4H2  = 3Fe  + 4H2O. 

In  the  first  instance  the  excess  of  water  acts.  Some  of  its  molecules  are  dissociated  ; 
oxygen  combines  with  iron,  while  the  liberated  hydrogen  is  carried  away  by  the  excess 
of  steam.  In  the  second  case,  we  can  suppose  that  some  of  the  hydrogen  molecules  are 
dissociated,  the  free  hydrogen  atoms  withdraw  oxygen  from  the  iron  oxide  and  form  water 
with  it,  which  is  removed  by  the  excess  of  hydrogen,  and  thus  prevented  from  acting  on 
the  reduced  iron. 

If,  however,  iron  and  steam  be  heated  in  an  enclosed  space  for  every  temperature 
above  150°,  where  action  begins,  there  will  occur  a state  of  equilibrium  : ferrous- ferric 
oxide,  iron,  water,  and  hydrogen  will  be  present  together  in  a ratio  which  does  not  alter 
for  any  definite  temperature,  i.  e.,  in  a definite  period  of  time  as  much  iron  will  be  pro- 
duced from  the  ferrous-ferric  oxide  by  the  hydrogen  as  is  oxidized  by  the  water  to  ferrous- 
ferric  oxide.  This  may  be  expressed  thus  : 

FegO,  + 4H2  3 Fe  -f  4H2O. 

(See  pp.  52,  68,  91.) 

Similarly,  hydrogen  chloride  is  decomposed  at  a red  heat  by  oxygen 
with  the  formation  of  steam  and  chlorine  gas,  while  in  turn  steam  and 
chlorine  gas  are  transposed  into  hydrogen  chloride  and  oxygen.  As 
hydrogen  chloride  is  more  stable  than  water  and  only  dissociated  at 
high  temperatures,  its  formation  is  the  predominating  process  at  a red 
heat.  Here,  however,  in  the  course  of  time,  a state  of  equilibrium 
appears  for  every  given  temperature  and  pressure. 

The  investigation  of  the  dependence  of  this  state  of  equilibrium  upon  temperature, 
pressure,  duration  of  action,  the  quantities,  concentration,  and  solubility-proportions  of  the 
reacting  substances,  has  become  of  the  greatest  importance  for  theoretical  chemistry,  for 
the  reason  that  here  for  the  first  time  chemical  phenomena  are  capable  of  a mathematical 
treatment.  The  fundamental  investigations,  both  theoretical  and  experimental,  made 
upon  the  law  of  mass-action,  which  will  not  be  entered  upon  here,  are  due  to  the  Nor- 
wegians C.  M.  Guldberg  and  P.  Waage  [Etudes  sur  les  affinites  chimiques,  Progr.  de 
I’Universite  Christiania,  1867.  See  also  Jr.  prakt.  Ch.  [2]  19  (1879),  69].  For  study 
use  the  works,  referred  to  on  p.  49,  by  Lothar  Meyer  ; also  the  Theoretical  Chemistry 
by  W.  Nernst,  2 Aufl.,  1898. 


THE  QUANTITATIVE  COMPOSITION  OF  WATER. 

The  weight  and  volume  relations  by  which  hydrogen  and  oxygen  com- 
bine to  form  water  constitute  the  most  important  basis  for  the  determina- 
tion of  the  ratio  of  the  atomic  weights  of  these  two  elements  to  each  other. 
To  determine  this  ratio  with  such  accuracy  that  no  question  can  longer 
exist  has  been  the  purpose  of  innumerable  investigations  since  the  time  of 
Dalton.  Numerous  difficulties  had  to  be  surmounted.  They  were  found 
particularly  when  seeking  to  obtain  the  gases  perfectly  dry,  in  weighing 


96 


INORGANIC  CIlEMIS'l'kV. 


ihcni,  in  measuring  them,  and  in  endeavoring  to  combine  them,  d’he 
judgment,  care,  and  untiring  j)atience  reviuisite  for  tlie  solution  of  tin’s  aj)- 
parently  simple  problem  may  be  gathered  from  the  magnificent  researches, 
already  referred  to,  of  the  American  Edward  W.  Morley  (pj).  44,  73). 

d'he  composition  of  water  has  been  chiefly  determined  by  its  synthesis 
which  can  be  followed  quantitatively  in  its  details,  as  both  oxygen  and 
hydrogen,  as  well  as  the  water  jiroduced  from  them,  are  weighed.  'J’he 
hydrogen  is  liberated  either  by  the  electrolysis  of  dilute  suliihuric  acid  or 
sodium  hydroxide  (Thomsen,  ]>.  42),  and  according  to  Reiser,  is  best 
weighed  as  jialladium  hydride  (p.  46).  'I'he  oxygen  is  obtained  by  heat- 
ing potassium  chlorate.  Chemists  have  frefiuently  been  satisfied,  instead 
of  weighing  the  three  bodies,  hydrogen,  oxygen,  and  water,  to  merely 
weigh  two  of  them.  This  was  true  of  Terzelius  and  Diilong  (1819),  Erd- 


Fig.  46. 


mann  and  Marchand  (1842),  and  Dumas  (1843),  who  proceeded  according 
to  a method  which  has  become  classic:  hydrogen  was  conducted  over 
ignited  copper  oxide,  which  was  reduced  to  metallic  copper,  with  the 
formation  of  water: 

CuO  + H2  = Cu  -f  H,0. 

Cupric  oxide.  Copper. 

Heat  a weighed  portion  of  cupric  oxide  (containing  a definite  amount 
of  oxygen)  in  a stream  of  ])ure,  dry  hydrogen,  and  weigh  the  quantity  of 
water  obtained.  The  operation  can  be  executed  in  the  apparatus  repre- 
sented in  Fig.  46.  It  was  em[)loyed  by  Dumas.  The  hydrogen  gener- 
ated in  the  flask  A is  washed  in  B,  and  then  dried  in  the  tubes  C,  D,  and 
/f,  which  contain  substances  that  will  absorb  water.  The  bulb  tube  F, 
of  difficultly  fusible  glass,  contains  a weighed  amount  of  cupric  oxide,  and 
is  heated  with  a lamp.  The  water  which  forms  collects  in  the  bulb  G, 
and  is  com])letely  absorbed  in  the  tube  H.  Hydrogen  is  led  over  the 
cupric  oxide  until  it  is  reduced  to  red  metallic  copper,  then  allowed  to 


THE  MOLECULAR  FORMULA  OF  WATER. 


97 


cool,  when  7^  is  weighed  alone  and  G and  iT"  together.  The  loss  in  weight 
of  expresses  the  quantity  of  oxygen  which  has  combined  with  hydrogen 
to  produce  water.  The  increase  in  weight  of  and  ZTgives  the  quantity 
of  water  that  was  formed.  The  difference  shows  the  amount  of  hydrogen 
in  water. 

Cooke  and  Richards  modified  this  method  in  that  they  conducted  a 
weighed  amount  of  hydrogen,  by  means  of  a dry  air  current,  over  the 
ignited  copper  oxide,  and  then  determined  the  weight  of  the  water 
which  resulted. 

The  composition  of  water  by  weight  can  also  be  ascertained  from  the 
ratio  of  the  specific  gravities  of  the  two  gases  and  the  volume  ratio 
according  to  which  they  combine.  The  most  recent  researches  (Scott, 
Leduc,  Morley)  indicate  that  2.0027  volumes  of  hydrogen  unite  with  i 
volume  of  oxygen  to  form  water.  As  one  liter  of  oxygen  weighs  1.4291 
gram  and  one  liter  of  hydrogen  0.08988  gram,  their  weight  ratio  in  water 
would  be 

I.4291  : 2.0027  • 0.08988  = 16  : 2.0153. 

Thus  we  ascertain  that  in  100  parts  of  water,  by  weight,  there  are 

1 1. 2 parts  Hydrogen 
88.8  “ Oxygen 

100.0  “ Water. 

Ostwald  in  his  Lehrbuch  der  allgemeinen  Chemie,  2 Aufl.  (1891),  i,  43,  gives  a 
review  of  the  history  of  these  important  investigations.  See  also  the  researches  of 
Morley  and  of  Thomsen  (p.  44). 


THE  MOLECULAR  FORMULA  OF  WATER.  ATOMIC  WEIGHTS  OF  HYDROGEN 

AND  OXYGEN. 

The  molecular  weight  of  water  according  to  Avogadro’s  law  is  18.02 
and  that  of  hydrogen  is  2.02  if  the  molecular  weight  of  oxygen  be  placed 
at  32.  As  to  the  number  of  atoms  combined  to  a molecule  we  can  con- 
clude from  what  has  been  previously  said  that  the  molecules  of  hydro- 
gen, oxygen,  chlorine  and  hydrogen  chloride  contain  at  least  two  atoms 
each  (p.  75).  The  facts  about  to  be  presented  make  it  probable  that 
these  molecules  do  not  consist  of  more  than  two  atoms  and  that  the 
molecular  formula  H^O  represents  water  (p.  73). 

One  volume  of  acpieous  vapor  contains  only  one-half  as  much  oxygen 
as  an  equal  volume  of  oxygen  gas,  and  in  an  equal  volume  of  any  gaseous 
oxygen  compound  less  oxygen  has  never  been  observed.  Hence  it  may 
fairly  be  concluded  that  there  is  only  one  atom  of  oxygen  present  in  the 
molecule  of  water,  and  that  the  oxygen  molecule  itself  consists  of  two  atoms. 

Again,  one  volume  of  aqueous  vapor  contains  just  as  much  hydrogen  as 
an  equal  volume  of  hydrogen  gas,  and  twice  as  much  as  hydrogen  chlo- 
ride gas.  Since,  therefore,  in  equal  volumes  of  other  gaseous  bodies  less 
hydrogen  and  less  chlorine,  than  in  hydrogen  chloride,  have  never  been 
observed,  it  maybe  assumed  that  the  molecule  of  hydrogen  chloride  con- 
tains one  atom  each  of  hydrogen  and  chlorine,  and  further  that  the 
9 


98 


INORGANIC  CHKMISTKY. 


hydrogen  molecule  consists  of  two  atoms  of  hydrogen,  the  water  molecule 
of  one  atom  of  oxygen  and  two  atoms  of  hydrogen.  In  this  manner 
we  arrive  at  the  molecular  and  atomic  weights  given  on  p.  79;  whereby 
the  atomic  weight  of  hydrogen  1.008  is  reduced  to  i.oi. 

It  seems  practicable,  for  considerations  like  those  given,  to  introduce  several  new 
terms.  As  the  sj)ecific  volume  Vs  of  a gas  we  have  (iesignated  that  volume  which  1 
gram  of  the  gas  would  occupy  under  normal  conditions  (p.  45).  d'he  quantity  of  a sub- 
stance expressed  by  the  molecular  weight  in  grams  slnndd  he  termed  the  g?'aiii-i?iolecule 
or  the  mol.  One  mol  of  oxygen  would  then  under  normal  conditions  he  32  grams.  'I’he 
volume  Vm,  which  would  he  occupied  by  i mol  of  a gas  of  the  molecular  weight  M under 
normal  conditions — the  violectdar  volmne — would  he  the  sj)ecific  volume  multiplied  by 
the  molecular  weight  : 

Vm  = Vs . M = 22.4  liters. 

The  mol-volume  is  the  same  for  all  gases : a new  expression  for  the  so-called  Avo- 
gadro  law.  It  equals  22.4  liters  ; i.  e.,  2.02  grams  of  hydrogen  70.9  grams  of 

chlorine  (Cl.^),  36.46  grams  of  hydrogen  chloride  (IlCl),  32  grams  of  oxygen  (02)>  i^-02 
grams  of  acjueous  vapor  occupy  under  normal  conditions  a volume  of  22.4  liters.  This 
is  only  approximately  correct  because  of  the  deviations  of  gases  from  the  re(|uirements  of 
their  laws.  What  has  been  said  in  the  preceding  lines  may  he  expressed  thus  : .Since  in 
22.4  liters  of  a homogeneous  gaseous  compound  never  less  than  i.oi  grams  of  hydrogen, 
35.45  grams  of  chlorine,  16  grams  of  oxygen,  etc.,  have  been  discovered,  it  may  be  con- 
cluded that  these  express  their  atomic  values. 

After  having  thus  derived  the  molecular  formula  of  water,  and  the  atomic 
weights  of  oxygen  and  of  hydrogen,  we  deduce  the  following  conclusions; 

I.  Sixteen  ytarts  by  weight  of  oxygen  occupy  the  same  volume  as  i.oi 

parts  by  weight  of  hydrogen;  since  16  partsof 
the  former  unite  with  2.02  parts  of  the  lat- 
ter in  the  production  of  water,  one  volume  of 
oxygen  must  combine  with  two  volumes  of 
hydrogen. 

2.  In  equal  volumes  we  have  an  equal 
number  of  molecules ; n molecules  of  oxygen 
(Oj)  unite  therefore  with  2n  molecules  of 
hydrogen  (H^) ; the  same  yield  2n  mole- 
cules of  water;  consequently  two  volumes  of 
aqueous  vapor : 

2nH2  -p  n02  = 2nH20. 

2 vols.  I vol.  2 vols. 

According  to  the  above,  two  7'o/umes  of  hy- 
drogen and  one  vohi77ie  of  oxygen  co7iden5e  i7i 
their  u7iio7i  to  tivo  vohwies  of  aqueous  vapor. 

'These  conclusions  are  confirmed  by  the 
following  exiteriments : 

I.  Wlien  water  is  decomposed  by  the  electric 
current  in  a voltameter,  or,  more  suitably,  in  Hof- 
mann’s apparatus  (Fig.  38,  j).  77b  it  will  be  fouiul 
that  the  volume  of  the  separated  hydrogen  is  double 
that  of  the  oxygen.  This  can  also  be  proved  syn- 
thetically. Introduce  I volume  of  oxygen  and  2 
volumes  of  hydrogen  into  a eudiometer  lube  fdlcd  with  mercury  (see  Air),  and  let  the 


HYDROGEN  PEROXIDE. 


99 


electric  spark  pass  through  the  mixture.  This  will  unite  the  two  gases,  a small  quantity 
of  water  forming  at  the  same  time  ; all  the  gas  has  disappeared,  and  the  tube  is  com- 
pletely filled  with  mercury.  In  place  of  the  eudiometer,  the  apparatus  (devised  by  Hof- 
mann) pictured  in  P'lg.  47  may  be  advantageously  employed  in  this  experiment  (and 
also  in  many  others).  It  consists  of  a U-shaped  glass  tube,  one  limb  of  which,  open 
above,  is  provided  below  with  an  exit  tube.  The  other  limb  really  represents  a eudi- 
ometer ; it  is  divided  into  cubic  centimeters.  It  has  two  platinum  wires  fused  into  the 
upper  end,  and  provided  with  a stop-cock  to  admit  and  let  out  the  gases  and  thus  test  them. 
P'ill  the  tube  to  the  stop-coek  with  mercury,  and  run  into  the  eudiometer  limb  i volume 
of  oxygen  and  2 volumes  of  hydrogen.  The  side  exit  tube  serves  to  run  out  the 
mercury  to  the  same  level  in  both  tubes,  so  that  the  gases  are  always  measured  under 
the  same  atmospheric  pressure,  and  thus  their  volumes  are  easily  compared. 

2.  To  determine  the  volume  of  the  resulting  water  existing  as  aqueous  vapor  it  is 
only  necessary,  after  the  explosion,  to  convert  it  by  heat  into  steam.  The  apparatus 
(Fig.  48)  will  answer  for  this  purpose.  This  is  essentially  the  same  as  that  pictured  in 
Fig.  47,  with  the  eudiometer  limb  closed  above  and  surrounded  by  a wider  tube. 
Through  the  latter  conduct  the  vapors  of  some  liquid  boiling  above  100°  C.  (toluene, 
xylene  or  aniline).  These,  then,  pass  through  the  envelope  B,  and  are  again  condensed 


Fig.  48. 

in  the  spiral  tube  C.  The  quantities  of  hydrogen  and  oxygen  used  are  heated  to  the  same 
temperature,  their  volume  noted,  the  explosion  produced,  and  the  volume  of  the  result- 
ing aqueous  vapor  determined.  From  this  it  is  Wnd  that  the  volume  of  hydrogen  is  ^ 
of  the  volume  of  the  gas  mixture  ; and  3 volumes  of  oxyhydrogen  gas  yield  2 volumes 
of  aqueous  vapor. 


2.  HYDROGEN  PEROXIDE. 

H2O2  = 34-02. 

In  addition  to  water,  oxygen  forms  another  compound  with  hydrogen, 
known  as  hydrogen  peroxide.  It  was  discovered  by  Thenard  in  1818. 
It  is  produced  by  the  action  of  dilute  acids  upon  certain  peroxides,  such 
as  those  of  sodium,  calcium  and  barium.  It  is  usually  obtained  by  the 
action  of  hydrochloric  acid  upon  barium  peroxide  : 

BaOj  -j-  2IICI  = RaCh  + HjOj. 

Barium  Barium 

peroxide.  chloride. 


lOO 


INORGANIC  CHEMISTRY. 


Barium  peroxide  made  to  a paste  with  a little  water  (better,  the  hydrate  — see  Barium) 
is  introduced  gradually,  in  small  (juantities,  into  cold  hydrcK'hloric  acid,  diluted  with  3 
volumes  of  water.  Hydrogen  peroxide  and  barium  chloride  result  ; both  are  soluble  in 
water.  To  remove  the  second  from  the  solution,  add  to  the  latter  a .solution  of  silver  sul- 
phate as  long  as  a precipitate  is  formed.  Two  insoluble  compounds,  barium  sulphate  and 
silver  chloride,  are  produced  by  this  reagent  : 

BaCl,  f Ag2SO,  = BaSO,  -f  2AgCl. 

Remove  the  precipitate  by  filtration  and  concentrate  the  aqueous  .solution.  It  now  con- 
tains only  hydrogen  peroxide. 

In  making  the  peroxide,  carbon  dioxide  may  be  allowed  to  act  on 
barium  peroxide  susitendcd  in  water  ; 

Ba02  CO2  + II2O  BaCOg  -f 

The  insoluble  barium  carbonate  is  filtered  off  and  the  filtrate  concen- 
trated. 

Hydrogen  jieroxide  is  most  practically  obtained  by  adding  moist 
barium  hydrated  peroxide  (see  llarium)  to  cold  dilute  sulphuric  acid,  d'he 
reaction  occurs  according  to  the  following  equation  : 

BaOj  + 112^^4  = BaSO^  -|-  II2O2. 

When  the  acid  is  almost  neutralized,  filter  the  solution,  and  from  the 
filtrate  carefully  precipitate  the  slight  quantity  of  free  sulphuric  acid  with 
a dilute  barium  hydroxide  solution,  then  concentrate  the  liquid.  Dry 
commercial  peroxide  of  barium  is  not  applicable  for  the  above.  A dilute 
solution  of  hydrogen  peroxide  is  very  readily  prepared,  if  sodium  peroxide 
(obtainable  by  fusing  sodium  in  the  air)  is  added  to  dilute  tartaric  acid. 

A 45  per  cent,  solution  can  be  readily  made  by  evaporating  the  dilute  aqueous  hydrogen 
peroxide  solutions  on  a water-bath  at  a temperature  not  exceeding  70°.  The  loss  will  be 
very  slight.  By  extracting  such  a solution  with  ether  and  allowing  the  latter  to  evaporate 
a more  concentrated  product  can  be  obtained  and  finally  an  almost  anhydrous  peroxide 
will  remain. 

The  simplest  means  of  preparing  very  pure  hydrogen  peroxide  (99.7  per 
cent.)  consists  in  distilling  the  aqueous  solution  (for  example,  the  commer- 
cial 3 per  cent,  solutions)  under  reduced  pressure.  At  84-85°,  under 
68  mm.  pressure,  or  at  69.2°  and  26  mm.  pressure,  almost  perfectly  anhy- 
drous hydrogen  peroxide  distils  over  [Wolffenstein,  Ber.  27  (1894),  ii, 
3307;  Spring,  Zeit.  f.  anorg.  Chem.  8 (1895),  424;  Briihl,  Ber.  28 

(1895),  2847-] 


Besides  these  decompositions  of  metallic  peroxides,  other  methods  exist  for  preparing 
hydrogen  peroxide  (in  small  quantity).  Thus,  it  ari.ses frequently  in  slow  oxidations,  when 
ozone  is  also  ])roduced.  If  phosphorus,  covered  with  water,  be  allowed  to  oxidize  in  the  air, 
hydrogen  peroxide  will  be  found  in  the  water,  and  the  surrounding  air  will  contain  ozone 
fp.  85).  Or,  if  a flask  filled  with  air  be  shaken  with  zinc  and  water  or  dilute  sulphuric 
acid,  hydrogen  peroxide  will  be  ])rodiiccd.  It  is  destroyed  again  by  the  prolonged  action  of 
the  zinc.  If  zinc  amalgam,  in  (he  presence  of  milk  of  lime,  be  .shaken  with  caustic  potash, 
calcium  ])eroxidc,  OaO,^,  will  be  at  once  ])recii)itated.  Copper,  lead,  iron,  and  other  heavy 
metals  do  the  same  when  agitated  with  more  or  less  dilute  sulphuric  acid,  and  we  find  the 


HYDROGEN  PEROXIDE. 


lOI 


same  result  by  the  oxidation  of  many  organic  substances,  e.  g. , pyrogallic  acid  and  tannin  on 
exposure  to  the  air.  In  combustions,  if  the  flame  be  cooled  suddenly,  we  have  formed,  very 
often,  slight  quantities  of  hydrogen  peroxide  (and  ozone),  e.  g.,  in  bringing  a hydrogen  or 
carbon  monoxide  flame  in  contact  with  water  (Traube,  Her.  26  (1893),  ii,  1471,  14761. 
The  explanation  offered  for  this  formation  of  hydrogen  peroxide  (and  ozone)  is,  that  in 
the  oxidations,  the  oxygen  molecules  are  torn  asunder,  and  the  nascent  oxygen  atoms 
oxidize  the  water  to  a slight  degree  to  hydrogen  peroxide,  and  oxygen  to  ozone.  The 
rare  occurrence  of  ozone  is  due  either  to  its  difficult  formation,  or  to  the  fact  that  it  is 
readily  decomposed  by  the  reacting  bodies  (zinc,  etc.).  This  is  also  the  case  wilh  hydro- 
gen peroxide.  The  appearance  of  hydrogen  peroxide  in  the  oxidation  of  plu)Sj)horus 
seems  to  prove  that  it  can  be  formed  by  the  oxidation  of  wate7\  This  seems  to  be  con- 
firmed by  its  production  on  shaking  turpentine  oil  with  water  and  air,  or  if  ozone  be  con- 
ducted into  ether,  and  the  ozonized  product  shaken  with  water.  It  appears  probable, 
however,  that  in  some  oxidation  reactions,  the  formation  of  the  hydrogen  peroxide  is  a 
consequence  of  the  reduction  of  oxygen  (Hoppe-Seyler  and  Traube).  It  may,  for  ex- 
ample, be  assumed  that  when  zinc  (lead,  iron)  is  shaken  with  air  and  water  (or  dilute 
sulphuric  acid),  the  latter  is  decomposed  in  such  a manner  that  the  hydroxyl  group  com- 
bines with  the  zinc  to  hydroxide,  and  the  liberated  hydrogen  then  yields  hydrogen 
peroxide  with  oxygen  : 

Zn  + 2OHH  + 02  = Zn(0H)2  + H2O2. 

Zn  + H2O2  = Zn  (011)2. 

A confirmation  of  this  supposition  is  found  in  the  electrolysis  of  water,  where  we  discover 
hydrogen  peroxide  appearing  at  the  negative  pole  (where  hydrogen  is  found)  if  air  or 
oxygen  be  conducted  through  the  solution,  2H  + 02=  H2O2.  It  is  verified,  too,  in 
the  production  of  hydrogen  peroxide  upon  shaking  palladium  hydride  with  water  and  air 
(Traube)  : 

Pd4H2  + 02  = 4Pd  + H2O2. 

The  excess  of  palladium  hydride  further  decomposes  the  peroxide  which  was  formed  : 

Pd,H2  + H2O2  = 4Pd  + 2H2O. 

In  all  these  examples  we  can  explain  the  formation  of  the  peroxide  by  the  action  of 
nascent  hydrogen  upon  oxygen.  It  is,  however,  not  true  that  hydrogen  peroxide  is 
formed  only  by  the  reduction  of  molecular  oxygen  (see  above).  An  evidence  of  this  is 
the  fact  that  hydrogen  peroxide  is  produced  at  the  anode  in  the  electrolysis  of  sulphuric 
acid  ; its  appearance  here  is  due  to  the  decomposition  of  the  persulphuric  acid  (H2S20g) 
(Richarz). 

The  production  of  hydrogen  peroxide  in  oxidations  has  led  to  the  assumption  that  all 
oxidations  are  conditioned  by  the  transitory  formation  of  hydrogen  peroxide — Traube’ s oxi- 
dation theory.  It  was  supposed  that  the  proof  of  this  could  be  found  in  the  circumstance 
that  various  oxidations,  e.  g.,  the  union  of  carbon  monoxide  with  oxygen  to  form  carbon 
dioxide,  could  only  occur  with  ease  in  the  presence  of  aqueous  vapor  (Dixon)  : CO  + 2OIIH 
-p  O2  = CO2  + II2O  + 112^^2-  More  careful  investigations  have,  however,  demonstrated 
that  the  presence  of  moisture  is  not  absolutely  e.ssential  in  oxidations.  Carbon  monoxide 
and  oxygen  also  combine  to  carbon  dioxide  when  perfectly  dry  if  the  temperature  be  suf- 
ficiently high.  Their  union  in  the  presence  of  moisture  is  due  solely  to  the  fact  that  the 
following  transpositions,  CO  + H2O  = CO2  + Ikj  and  2II2  + O2  = 2H2O,  take  place 
more  readily  and  at  a lower  temperature  than  the  direct  union  of  carbon  monoxide  with 
oxygen  : 2CO  + 02  = 2CO2  (Lothar  Meyer ; see  p.  92). 


Hydrogen  peroxide,  concentrated  as  much  as  possible  under  diminished 
pressure,  is  a colorless,  syrupy  liquid  which  is  blue  in  color  in  a thick  layer. 
It  does  not  congeal  as  readily  as  water.  Its  specific  gravity  at  15°,  referred 
to  water  at  the  same  temperature,  is  1.49.  It  vaporizes  on  exposure  to  the 
air.  It  produces  a burning  sensation  and  white  spots  when  applied  to  the 


102 


INORGANIC  CHEMISTRY. 


skin,  ll  sometimes  cx])lo(ies  spontaneously  with  great  violence.  Its 
aqueous,  rather  concentrated  solutions  react  acid  and  have  a bitter, 
astringent  taste.  Hydrogen  peroxide  is  not  nearly  so  sensitive  to  heat 
as  was  tormerly  siq)posed,  provided  that  all  alkaline-reacting  compounds 
and  chemically  active  solids  are  absent.  It  is  miscible  in  all  proportions 
with  water.  Mixtures  of  the  conqiosition  H./).^  -f  H/)  and  -j-  2H2O 
solidify  below  — 20°.  Upon  warming  solutions  containing  more  than 
40  per  cent,  of  hydrogen  j)eroxide,  the  latter  volatilizes  in  large  amounts 
with  the  aqueous  vapor  and  for  the  most  part  without  decomposition 
(Wolffenstein,  p.  100). 

In  concentrated  solutions,  it  is  very  unstable,  and  easily  decomposed 
with  liberation  of  oxygen  ; in  more  dilute  acidulated  solutions  it  may  be 
preserved  longer.  According  to  Spring,  substances  such  as  ether  and 
alcohol,  which  reduce  the  surface  tension  of  the  liipiid,  tend  to  make  its 
solution  more  durable,  whereas  bodies  like  caustic  potash,  etc.,  which 
increase  the  surface  tension,  hasten  the  breaking  down  of  the  hydrogen 
peroxide.  In  consequence  of  its  ready  decomposition,  hydrogen  peroxide 
oxidizes  powerfully,  since  oxygen  appears statu  nascendi  {\).  53).  It 
converts  selenium,  chromium,  and  arsenic  into  their  corresponding  acids; 
sulphides  are  changed  to  sulphates  (PbS  to  PbSOi) ; from  lead  acetate 
solutions  the  peroxide  is  precipitated,  but  is  again  decolorized  by  the 
excess  of  hydrogen  peroxide.  Organic  dyestuffs  are  decolorized  and 
decomposed.  From  hydrogen  sulphide,  sulphur,  from  hydrogen  chloride 
and  iodide,  chlorine  and  iodine  are  set  free : 

H2O2  + 2HI  = 2H2O  + Ij. 

Sulphurous  acid  is  oxidized  to  sulphuric  acid  : 

H2SO3  + 11,0,  = H^SO,  + H^O. 

Thus  hydrogen  peroxide  acts  in  a manner  analogous  to  ozone ; in  both 
there  exists  a loosely  combined  atom  of  oxygen,  which  can  readily  be 
transferred  to  other  bodies. 

Hydrogen  peroxide  acts  very  slowly  upon  a neutral  potassium  iodide 
solution,  while  ozone  separates  iodine  at  once  ; but  if  platinum-black,  fer- 
rous sulphate,  or  blood-corpuscles  (see  p.  88),  be  added  to  the  solution, 
iodine  immediately  separates,  and  colors  added  starch-paste  a deep 
blue. 

In  all  these  cases  the  action  of  hydrogen  peroxide  is  oxidizing.  Some 
substances,  on  the  other  hand,  are  reduced  by  it,  oxygen  separating  at 
the  same  time ; this  is  true  of  certain  unstable  oxides,  peroxides,  and  the 
highest  oxidations  of  some  metals,  like  MuaO;,  and  CrO.,.  Thus,  argen- 
tic, mercuric,  and  gold  oxides  are  reduced  to  a metallic  state  with  an 
energetic  evolution  of  oxygen  : 

Ag.p  4 II.A  = 2Ag  + II2O  -f  O2. 

Lead  peroxide  is  changed  to  lead  oxide : 

PhO^  -I-  U,0,  = PbO  -f  1^0  -p  O,. 

In  the  jiresence  of  acids,  the  solution  of  potassium  permanganate  is 


REACTIONS  FOR  THE  DETECTION  OF  HYDROGEN  PEROXIDE.  IO3 

decolorized  and  changed  to  a manganous  salt  (see  below).  In  the  same 
way  chromic  acid  and  its  salts  are  altered  to  chromic  oxide : 

2Cr03  -f-  3H2O2  = Cr203  + 3H2O  + 3O2. 

Ozone  and  hydrogen  peroxide  gradually  decompose  into  water  and 
oxygen  : 

O3  + 11.^02  = 02  + H2O  Oj. 

Chlorine  in  aqueous  solution  is  oxidized  to  hypochlorous  acid  by 
hydrogen  peroxide,  CI2  -f-  H2O2  = 2HOCI,  but  is  again  reduced  by  an 
excess  of  the  latter  : 

HCIO  + H2O2  = HCl  + H,0  -f  O2. 

Finally,  hydrogen  peroxide  may  be  decomposed  into  water  and  oxygen 
by  many  bodies,  especially  when  the  latter  exist  in  a divided  condition  ; 
they  are  not  in  the  least  altered.  Gold,  platinum,  silver,  manganese  per- 
oxide, and  carbon  act  in  this  way.  Such  reactions,  in  which  fhe 
reacting  substances  undergo  no  perceptible  changes,  are  designated  cata- 
lytic (xaraXooj,  I open)  (compare  p.  102).  In  many  cases  these  may  be 
explained  by  the  previous  formation  of  intermediate  products,  which 
subsequently  react  upon  one  another.  Thus,  we  can  suppose  that  in  the 
action  of  silver  and  gold  upon  hydrogen  peroxide  oxides  first  result,  but 
these  are  afterward  reduced  by  it  in  the  manner  mentioned  above  (see 
p.  87). 

REACTIONS  FOR  THE  DETECTION  OF  HYDROGEN  PEROXIDE. 

Hydrogen  peroxide  decomposes  potassium  iodide  very  slowly  ; in  the  presence  of  ferrous 
sulphate,  however,  iodine  separates  at  once,  and  is  recognized  by  the  blue  color  it  yields 
with  starch-paste.  In  the  same  way  guaiacum  tincture,  in  the  presence  of  ferrous  sul- 
phate, is  immediately  colored  blue,  and  an  indigo  solution  is  decolorized.  The  most 
characteristic  test  for  the  peroxide  is  the  following  : introduce  hydrogen  peroxide  into  a 
chromic  acid  solution,  add  a little  ether  and  shake  thoroughly  ; ths  supernatant  ethereal 
layer  will  be  colored  blue  (compare  Chromic  Acid). 

A solution  of  titanic  acid  in  sulphuric  acid  (diluted  strongly  with  water),  is  also  a deli- 
cate reagent  ; it  gives  an  orange-yellow  color  with  traces  of  hydrogen  peroxide.  This 
reaction  can  also  be  applied  in  the  presence  of  persulphuric  acid. 

Hydrogen  peroxide  is  determined  quantitatively  by  oxidation  with  potassium  perman- 
ganate (see  Manganese).  The  latter  is  added  to  the  solution,  acidified  with  sulphuric 
acid,  until  a permanent  coloration  occurs.  The  reaction  proceeds  according  to  the 
equation : 

2KMn04  -|-  3H2SO4  -j-  5H2O2  = 2MnS04  -f-  K2SO4  -f-  SH^O  -(-  5O2. 

Or  the  liquid  to  be  examined  (rain-water)  for  hydrogen  peroxide  is  shaken  in  a stop- 
pered glass  with  a 5 per  cent,  solution  of  pota.ssium  iodide  and  some  starch-paste, 
allowed  to  stand  several  hours,  and  the  iodine  which  separates  is  then  determined  color- 
irnetrically  (Schone). 

Thermo-chemical  Deportment. — The  great  reactivity  ef  hydrogen  per- 
oxide, its  various  modes  of  formation  and  its  transpositions  are  fully 
exi)lained  by  its  thermal  relations.  Compared  with  water,  it,  like 
ozone,  is  an  endothermic  compound,  i.  e.,  it  contains  more  energy  than 
water : (H2G,0)  = — 23  Cal.  Therefore,  its  formation  from  the  latter 
requires  the  addition  of  energy.  Its  ])roduction  by  the  oxidation  of 


104 


INORGANIC  CHEMISTRY. 


water  is  excei)lional,  and  occurs  with  (lifficulty.  It  loses  energy  (heat) 
and  readily  changes  to  more  stal)le  water.  Ij'ke  ozone,  its  oxidations 
proceed  more  energetically  than  those  with  free  oxygen,  because  23 
Cal.  more  are  disengaged. 

The  production  of  hydrogen  peroxide  by  the  transposition  of  barium 
peroxide  and  hydrochloric  acid  proceeds  with  the  liberation  of  heat: 

BaO^  4-  2lICl,Aq  = P>aCl2,Aq  4 H2O2  . . . 4-22.0  Cal. 

Hydrogen  peroxide  is  similarly  formed  from  other  peroxides,  e.  g., 
potassium,  calcium,  and  zinc  i)eroxides.  The  superoxides  or  dioxides 
of  manganese  and  lead  (Midland  1*1)02)  not  yield  h\drogen  i)erox- 
ide  with  acids.  This  is  due  to  the  fact  that  they  are  differently  con- 
stituted chemically  from  the  other  j)eroxides. 

Hydrogen  peroxide  occurs  in  slight  (juantity  in  the  air  and  is  detected 
in  almost  all  rain-water  and  in  snow — but  not  in  natural  dew  and  frost. 
Its  quantity  varies  from  0.5  to  i milligram  in  a liter  of  rain.  Its  forma- 
tion in  the  air  is  probably  induced  by  the  action  of  ozone  upon  ammonia, 
whereby  ammonium  nitrite,  hydrogen  peroxide,  and  oxygen  result  (Carius). 

Analysis  shows  that  hydrogen  |)eroxide  consists  of  i.oi  parts  of  hydrogen 
and  16  parts  of  oxygen;  its  simj)lest  formula  would  therefore  be  HO. 
The  difficult  volatility  of  the  comj)ound,  and  the  reactions  already 
described,  cause  us  to  believe  that  the  molecule  of  hydrogen  ])eroxide 
is  more  complicated,  and  is  expressed  by  H2O2.  It  is  supposed  that  the 
peroxide  is  composed  of  two  groups  of  OH,  called  hydroxyl ; these  are 
combined  with  each  other  : HO  — OH. 


2.  SULPHUR. 

Atom  : S = 32.06.  Molecule  : S2  ==  64.12  (above  1000°  C.). 

Sulphur  is  distributed  throughout  nature,  both  free  and  in  a combined 
state.  In  volcanic  regions,  like  Sicily,  it  occurs  free,  and  there  it  forms 
vast  deposits,  mixed  with  gy])sum,  calcite  and  marl.  Its  compounds 
with  the  metals  are  known  as  blendes  or  glances.  In  combination  with 
oxygen  and  calcium  it  forms  calcium  sulphate,  the  widely  distributed  gyp- 
sum. It  is  also  ])resent  in  many  substances  of  the  vegetable  and  animal 
kingdoms — e.  g.,  in  the  albuminoid  bodies.  It  is  interesting  to  note  that 
many  bacteriae  and  algae  contain  as  much  as  one-fourth  of  their  weight 
of  sul})hur. 

d'o  obtain  sulphur,  the  natural  product  in  Sicily  is  arranged  in  heaps, 
covered  with  earth,  and  then  melted,  or  it  is  distilled  from  earthen 
retorts.  To  further  ])urify  this  crude  coanmercial  product  it  is  redistilled 
from  cast-iron  retorts,  and  when  in  a molten  condition  is  run  into  cylin- 
drical forms — stick  sulphur.  If  the  suli)hur  vapors  are  rapidly  cooled 
during  distillation  (which  occurs  by  conducting  them  into  a stone 
chaml)er  through  which  cold  air  circulates),  they  condense  to  a fine 
yellow  jiowder,  known  as  jlowers  of  sulphur  (Flores  sulphuris). 


SULPHUR.  105 

Sulphur  may  be  obtained  by  heating  the  well-known  pyrites  (FeS^) 
away  from  air  contact. 

Appreciable  quantities  of  sulphur  are  obtained  from  the  material  used  in  purifying  gas — 
Laming’ s substance  (ferric  hydrate,  lime  and  sawdust ; which  finally  contains  iron  sulphide 
and  sulphur),  as  well  as  in  the  LeBlanc  soda  process  (see  Soda). 

Free  sulphur  exists  in  several  allotropic  modifications  (see  p.  87). 

1.  Ordinary  octahedral  or  rhombic  sulphur  in  nature  in  beauti- 

ful, well-crystallized  rhombic  octahedra  (p.  35).  It  is  pale  yellow,  hard  and 
very  brittle;  on  rubbing,  it  becomes  negatively  electrified.  The  specific 
gravity  of  this  variety  equals  2.07.  It  dissolves  with  difficulty  in  alcohol 
and  ether;  but  is  more  readily  soluble  in  hydrocarbons  and  ethereal  oils. 
The  best  solvents  are  sulphur  monochloride  (S2CI2)  and  carbon  bisulphide 
(CS2)  ; 100  parts  of  the  latter  at  22°  dissolve  46  parts  of  sulphur.  By 
slow  evaporation  of  the  solutions  sulphur  crystallizes  in  transparent,  lus- 
trous, rhombic  octahedra,  like  those  occurring  in  nature.  It  fuses 
at  114.5°  C.  to  a yellow,  mobile  liquid,  which  upon  further  heating  be- 
comes dark  and  thick,  and  at  250°  is  so  viscid  that  it  cannot  be  poured 
from  the  vessel  containing  it.  Above  300°  it  again  becomes  a thin 
liquid,  boils  at  448°,  and  is  converted  into  an  orange-yellow  vapor. 

2.  The  prismatic  or  77ionoclinic  sulphur  is  obtained  from  the  rhombic 
when  the  latter  is  heated  to  its  point  of  fusion ; on  cooling,  it  generally 
assumes  the  monoclinic  form  (rhombic  crystals  separate  at  90°  from 
sulphur  which  has  been  heated  beyond  the  point  of  fusion).  The  mono- 
clinic crystals  are  best  obtained  as  follows:  Fuse  sulphur  in  a clay  cruci- 
ble, allow  it  to  cool  slowly  until  a crust  appears  on  the  surface ; break 
this  open  near  the  side  and  pour  out  the  liquid- portion.  The  walls  of 
the  crucible  will  be  covered  with  long,  somewhat  curved,  transparent, 
brownish-yellow  needles,  or  prisms  of  the  monoclinic  system.  The  same 
are  obtained  when  a solution  of  sulphur  in  carbon  bisulphide  is  heated  to 
100°,  in  a sealed  tube,  and  then  gradually  allowed  to  cool;  mono- 
clinic crystals  at  first  separate,  and  later,  at  lower  temperatures,  rhombic 
octahedra.  The  monoclinic  crystals  separated  from  the  solution  are 
almost  colorless  and  perfectly  transi)arent. 

Prismatic  or  octahedral  crystals  may  be  obtained  from  a supersaturated  benzene 
solution  of  sulphur,  by  adding  small  fragments  of  the  corresponding  crystals  to  the 
solution. 

This  form  of  sulphur  has  a lower  specific  gravity  (=1.96)  and  fuses 
above  1 20°.  It  is  soluble  in  the  same  solvents  as  the  rhombic  variety.  It 
is  very  unstable;  the  transparent  prisms  and  needles  become  opaque  and 
pale  yellow  at  ordinary  temperatures,  and  specifically  heavier,  and  pass 
over  into  an  aggregate  of  rhombic  octahedra  retaining  the  external  pris- 
matic form.  Stick  sulphur  deports  itself  similarly;  the  freshly  moulded 
sticks  are  comjiosed  of  monoclinic  prisms,  but  in  time  their  specific 
gravity  changes  and  they  are  converted  into  the  rhombic  modification. 

3.  So/t,  plastic  sulphur  appesivs  to  consist  of  two  modifications.  It  is 
obtained  when  sulphur  heated  above  230°  is  poured  in  a thin  stream 
into  water ; it  then  forms  a soft,  fusible  mass,  of  a yellowish-brown  color. 


INORGANIC  CHEMISTRY. 


lOG 

and  its  specific  gravity  ccpials  1.92.  In  a few  days  it  hardens,  and  is 
converted  into  the  rhombic  variety.  At  95°  the  conversion  is  instan- 
taneous and  accoini)anied  by  the  evolution  of  considerable  heat.  It  is 
only  partly  soluble  in  carbon  bisulphide,  leaving  an  amorphous  powder 
undissolved — amorphous  insoluble  sulphur.  As  it  reverts  to  the  rhombic 
at  100°  molten  suljdiur  must  be  quickly  chilled  in  order  to  obtain  much 
of  it.  It  is  also  ])roduced  when  light  acts  upon  dissolved  or  fused  sul- 
phur, and  in  the  decomposition  of  the  halogen-suljjhur  compounds  by 
water.  Flowers  of  sulphur,  obtained  by  cooling  sulphur  vai)or  quickly, 
are  for  the  most  j)art  insoluble  in  carbon  bisulphide. 

On  adding  hydrochloric  acid  to  polysulphide  solutions  of  potassium  or 
calcium,  suli)hur  sei)arates  as  a fine,  white  powder,  known  as  7fiilk  of  sulphur 
(Lac  sulphuris)  : 

K2S5  + 2IICI  2KCI  -f  II^S  -f  4S. 

This  is  amorphous,  soluble  in  carbon  bisulphide,  and  gradually  passes 
into  the  rhombic  form. 

The  existence  of  these  various  modifications  of  sulphur,  like  that  of  ordi- 
nary oxygen  and  ozone,  may  be  attributed  to  the  fact  that  the  sulphur 
molecules  do  not,  under  all  circumstances,  consist  of  the  same  number 
of  atoms.  This  supposition  is  confirmed  by  the  deportment  of  sulphur 
vapor.  The  density  of  the  latter  at  500°  has  been  found  to  equal 
192  (O2  = 32).  The  vapor  density  steadily  diminishes  with  increase  of 
temperature  and  becomes  constant  at  1000°,  and  equals  64.  Since 
the  atomic  weight  of  sulphur  equals  32,  it  follows  that  at  1000°  its 
molecules  consist  of  two  atoms  (S2  = 64  = 32  X 2).  At  lower  temperatures 
the  sulphur  vapors  appear  to  contain  molecules  consisting  of  more  than 
two  atoms;  thus  at  500°,  where  the  vapor  density  equals  192,  the  molecule 
consists  of  six  atoms.  According  to  this  the  hexatomic  sulphur  molecules 
dissociate,  on  further  heating,  and  break  down  into  normal  diatomic  mole- 
cules ; the  dissociation  begins  at  700°  and  is  complete  at  1000°  (see  p.  93). 
Since,  therefore,  the  sulphur  molecules  in  vapor  form  consist  of  two 
atoms  at  very  high  temperatures  and  of  more  atoms  at  lower,  we  may 
assume  that  the  molecules  in  the  liquid  and  solid  condition  are  more 
complicated,  and  that  the  various  allotropic  modifications  are  influenced 
by  the  number  of  atoms  contained  in  the  molecules.  Other  solid  metal- 
loids— e.  g.,  selenium,  phosphorus,  arsenic,  carbon,  and  silicon — occur  in 
different  modifications.  As  yet  we  have  no  means  of  ascertaining  the 
molecular  size  of  the  elements  in  liquid  and  solid  conditions;  there  is 
much,  however,  favoring  the  idea  that  when  free  they  consist  of  complex 
atomic  groups.  Another  explanation  of  allotropy  will  be  given  in  con- 
nection with  the  different  varieties  of  ])hosphorus. 


Chemical  J^roperlies. — In  its  chemical  behavior  sulphur  is  very  similar 
to  oxygen.  It  unites  directly  with  most  of  the  elements.  It  ignites  and 
burns  with  a pale  bluish  flame,  forming  sulphur  dioxide  (SO.^)  when 
heated  to  260°  in  the  air.  This  union  with  oxygen  occurs  gradually 


HYDROGEN  SULPHIDE. 


107 


even  at  lower  temperatures  (about  180°);  in  the  dark  it  is  accompanied 
by  a white  phosphorescent  flame.  Nearly  all  the  metals  combine  with  it 
to  form  sulphides.  By  rubbing  mercury,  flowers  of  sulphur  and  water 
together,  we  obtain  black  mercury  sulphide.  A moist  mixture  of  iron 
filings  and  sulphur  glows  after  a time.  Copper  and  iron  burn  in  sulphur 
vapor. 

The  sulphides  are  analogous  to  the  oxides,  exhibit  similar  reactions, 
and  in  the  main  possess  a similar  composition,  as  may  be  seen  from  the 
following  formulas: 


H.^O,  Water. 

KOH,  Potassium  hydrate. 
BaO,  Barium  oxide. 

COj,  Carbon  dioxide. 

K.^COj,  Potassium  carbonate. 


HgS,  Hydrogen  sulphide. 

KSH,  Potassium  sulphydrate. 

BaS,  Barium  sulphide. 

CS.^,  Carbon  bisulphide. 

KjCSg,  Potassium  sulphocarbonate. 


COMPOUNDS  OF  SULPHUR  WITH  HYDROGEN. 

1.  HYDROGEN  SULPHIDE. 

H2S  = 34.08. 

In  nature  hydrogen  sulphide  occurs  principally  in  volcanic  gases  and 
in  the  so-called  sulphur  waters  (p.  91).  It  is  always  produced  in  the 
decomposition  of  organic  substances  containing  sulphur,  particularly  the 
albuminoids  by  which  the  sulphates,  e.  g.,  gypsum,  are  reduced  with  the 
formation  of  hydrogen  sulphide.  Hence  the  occurrence  of  the  gas  in 
eggs,  sewers,  etc.  It  may  be  formed  directly  from  its  constituents, 
although  in  small  quantity,  if  hydrogen  gas  be  conducted  through  boil- 
ing sulphur,  or  if  sulphur  vapors,  together  with  hydrogen,  be  conducted 
over  porous  substances  (pumice-stone,  bricks)  heated  to  500°.  Many 
sulphides  are  reduced  upon  ignition  in  a stream  of  hydrogen,  with  sepa- 
ration of  hydrogen  sulphide : 

Ag^S  + H,  = 2Ag  + H,S. 

For  its  production  acids  are  allowed  to  act  upon  sulphides.  Ordi- 
narily iron  sulphide  and  diluted  sulphuric  acid  are  employed ; the  action 
occurs  at  ordinary  temperatures  : 

FeS  + H.SO^  = FeSO,  -f  H^S. 

Iron  Sulphuric  Ferrous 
sulphide.  acid.  sulphate. 

The  operation  is  performed  either  in  a Kipp  apparatus  (p.  42)  or  in  the 
one  pictured  on  p.  43.  Hydrogen  sulphide  thus  obtained  contains  ad- 
mixed hydrogen,  in  consequence  of  metallic  iron  existing  in  the  sulphide. 
The  pure  gas  is  obtained  by  heating  antimony  sulphide  with  concentrated 
hydrochloric  acid  : 


Sb.Sj  -f  6IIC1  = 2SbCl3  -P  3lI.,S. 


io8 


INORGANIC  CHEMISTRY. 


Properties. — Hydrogen  sulphide  is  a colorless  gas,  having  an  odor  sim- 
ilar to  that  of  rotten  eggs;  inhaled  in  large  (juantities  it  has  a stupefying 
effect,  and  is  very  ])(hsonoiis.  At  medium  temperatures  it  eondenses 
under  a i)ressure  of  ry  atmospheres  (under  ordinary  pressure  at  — 74°)  to 
a colorless  licpiid  of  specific  gravity  0.9,  which  boils  at — 63.5°  under  760 
mm.  pressure,  and  at  — 91°  solidifies  to  a white  crystalline  mass.  It  is 
1,18  times  heavier  than  air.  One  volume  of  water  dissolves  3 or  4 times 
its  volume  of  the  gas  ; the  solution  jiossesses  all  the  jiroperties  of  gaseous 
hydrogen  suli)hide  and  is  therefore  called  hydrogen  sulphide  water. 

Ignited  in  the  air  the  gas  burns  with  a blue  flame,  water  and  suliihur 
dioxide  resulting : 

II,S  -f  3O  = 11,0  + SO,. 

With  insufficient  air  access,  or  when  the  flame  is  cooled  by  the  intro- 
duction of  a cold  body,  only  hydrogen  burns  and  sulphur  sejiarates  out  in 
a free  condition.  This  behavior  is  utilized  in  the  technical  preparation 
of  sulphur  from  soda  residues.  In  acpieous  solution  hydrogen  sulphide  is 
similarly  decomposed  by  the  oxygen  of  the  air  at  ordinary  temperatures, 
sulphur  sejiarating  as  a fine  powder : 

II,S  + O =r  H,0  + S. 

The  halogens  behave  like  oxygen ; the  hydrides  of  the  halogens  are 
formed  with  separation  of  sulphur : 

H,S  + I2  = 2HI  4-  S. 

This  reaction,  which  occurs  only  in  the  presence  of  water,  serves  for  the 
production  of  hydrogen  iodide  (p.  63). 

As  hydrogen  sulphide  has  a great  affinity  for  oxygen,  it  withdraws  the 
latter  from  many  of  its  compounds,  and  it  therefore  acts  as  a reducing 
agent.  Thus  chromic,  manganic,  and  nitric  acids  are  reduced  to  lower 
stages  of  oxidation.  On  pouring  fuming  nitric  acid  into  a vessel  filled 
with  the  dry  gas,  the  mixture  will  unite  with  a slight  explosion. 

Hydrogen  sulphide  possesses  weak  acid  properties,  reddens  blue  litmus- 
])aper,  forms  salt-like  compounds  with  bases,  and  is,  therefore,  termed 
hydrosulphuric  acid.  Nearly  all  the  metals  liberate  hydrogen  from  it, 
yielding  metallic  sulphides: 

Pb  -f  11,5  = PbS  + II,. 

Wi  h the  oxides  and  hydroxides  of  the  metals  hydrogen  sulphide  yields 
sulphides  and  suli)hydrates : 

KOI  I + II,S  = KSH  -p  11,0. 

Potassium 

hydrosulphide. 

PbO  + II,S  = PbS  -f  11,0. 

Lead 

sulphide. 

Sulphides,  therefore,  like  the  compounds  of  the  halogens  with  the 
metals,  may  be  viewed  as  the  salts  of  liydrosuliihuric  acid.  The  sulphides 
of  almost  all  the  heavy  metals  are  insoluble  in  water  and  dilute  acids; 


MOLECULAR  FORMULA  OF  HYDROGEN  SULPHIDE.  I09 

therefore,  they  are  precipitated  by  hydrogen  sulphide  from  solutions  of 
metallic  salts: 

CuSO,  + H^S  = CuS  4-  H^SO,. 

Copper  Copper  Sulphuric 

sulphate.  sulphide.  acid. 

The  precipitates  thus  obtained  are  variously  colored  (copper  sulphide, 
black;  cadmium  sulphide,  yellow;  antimony  sulphide,  orange),  and 
answer  for  the  characterization  and  recognition  of  the  corresponding 
metals.  Pai)er  saturated  with  a lead  salt  solution  is  at  once  blackened  by 
hydrogen  sulphide,  lead  sulphide  being  formed— a delicate  test  for  the 
gas. 

Thermo-chemical  Deportment. — Hydrogen  sulphide  is  a feebly  exother- 
mic compound.  When  hydrogen  gas  unites  with  amorphous  sulphur  to 
form  hydrogen  sulphide  4.7  Cal.  are  developed.  When  the  gas  dissolves 
in  much  water  its  heat  of  solution  equals  -{-4.6  Cal.,  so  that  the  total 
heat  of  formation  of  hydrogen  sulphide  in  dilute  aqueous  solution  is 
9.3  Cal. : 

(H2,S  — gas)  = 4.7  ; (H,S,Aq)  ==  4.6  ; (H,,S,Aq)  = 9.3. 

It  is  because  of  this  low  heat  of  formation  that  the  gas  is  produced  with 
such  difficulty  from  its  elements,  and  it  is  for  this  reason  that  it  is  so  readily 
dissociated  by  heat  into  its  elements.  Its  entire  chemical  deportment  is 
also  accounted  for  by  its  heat  of  formation  (p.  114). 


MOLECULAR  FORMULA  OF  HYDROGEN  SULPHIDE.  ATOMIC  WEIGHT  OF 

SULPHUR. 

Considerations  similar  to  those  which  led  us  to  adopt  the  molecular  formula  H2O  for 
water  impel  us  to  accept  the  formula  H.^S  for  hydrogen  sulphide.  Its  gas  density,  its 
analysis,  and  the  fact  that  in  an  equal  volume  of  any  sulphur-containing  gaseous  com- 
pound there  has  never  been  observed  less  sulphur  than  is  contained  in  hydrogen  sulphide 
speak  in  favor  of  the  assumption.  The  atomic  value  of  sulphur  is  therefore  ^2.06  (com- 
pare p.  97). 

From  the  molecular  formula  H^S  we  further  conclude  that  the  hydrogen  contained  in 
one  volume  of  hydrogen  sulphide  would  occupy  in  a free  condition  the  same  volume  as 
the  latter  : 

nll.^S  contains  nH2. 

I vol.  I vol. 

This  conclusion  is  verified  experimentally  as  follows  : In  a bent  glass  tube  filled  with  mer- 
cury (p.  78,  Fig.  41),  introduce  dry  hydrogen  sulphide  gas  ; then  in  the  bent  portion  place 
a piece  of  tin,  which  is  heated  by  a lamp.  The  sulphur  of  the  hydrogen  sulphide  combines 
with  the  metal  to  form  solid  tin  sulphide,  while  hydrogen  is  set  free  ; its  volume  is 
exactly  equal  to  the  volume  of  the  employed  hydrogen  sulphide.  The  quantity  of  sul- 
phur, 32  parts,  in  vapor  form,  at  1000°,  when  the  density  is  64  (p.  106)  combined 
with  hydrogen  (2  parts)  would  equal  exactly  half  the  volume  of  the  hydrogen  ; at  500°, 
however,  when  the  vapor  density  is  three  times  as  great,  it  will  equal  one-sixth  of  the 
volume  of  the  hydrogen.  Written  molecularly,  we  have  : 


At  500°  : Sg 

+ 

6112  = 

6II2S. 

I vol. 

6 vols. 

6 vols. 

At  1000°,  however  : S2 

2II2  = 

2II2S. 

I vol. 

2 vols. 

2 vols. 

I TO 


INORGANIC  CHEMISTRY. 


HYDROGEN  PERSULPHIDE. 

Just  as  hydrogen  ijeroxide,  is  formed  by  the  action  of  acids 

upon  some  peroxide,  so  may  hydrogen  persul])hide  lie  obtained  from 
metallic  persulphides  or  polysulphides.  Sodium  and  potassium  each 
form  four  iiolysulphides  which  are  available  for  this  jiurpose.  With 
sodium  for  exam])le  there  are,  in  addition  to  the  ordinary  suljjhide  Na^S, 
also  the  compounds  Na,^S^,  Na^S.,,  Na,^S^,  Na^S^.  J'\;r  (piite  a while  it  was 
thought  that  these  polysulphides  were  decomposed  by  dilute  acids  simi- 
larly to  the  decomposition  of  barium  jieroxide  by  dilute  sulphuric  acid  : 

V,vi(\  + 1 1, SO,  = P,aSO,  -f 

that  therefore  each  polysuljihide  had  a corresponding  hydrogen  polysul- 
phide. This  idea  was  enforced  by  the  fact  that  the  alkaloid  strychnine 
formed  a crystalline  product  with  H^S.,,  and  brucine  one  with 
Rebs,  however,  demonstrated  that  when  the  different  polysulphides  are 
decomposed  it  is  probably  always  hydrogen  pentasulphide,  H^S^,  which  is 
produced.  It  is  a light-yellow  colored,  transjiarent  mobile  oil  with  a 
peculiar  odor.  Its  sjiecinc  gravity  equals  1.7 1 at  15°.  When  dry  and 
away  from  air  contact  it  decomposes  slowly.  Water  resolves  it  quickly 
into  hydrogen  sulphide  and  sulphur,  which  separates.  Heat  induces  the 
same  decomposition  : 

H2S5  = S,  + H,S. 

Rebs  (see  Ann.  Chem.  (1888)  246,  354)  explains  the  formation  of  hydrogen  penta- 
sulphide  from  the  alkaline  bi-,  tri-,  and  tetrasulphides  in  the  following  equations: 

4Na2S,  + 8HC1  8NaCl  + 4H2S, ; 4H2S,  = + H2S. 

4Na2S3  + 8HC1  = 8NaCl  + 4H2S3 ; 4H2S3  = 2II2S5  -f  2H2S. 

4Na2S2  + 8HC1  = 8NaCl  + 4H2S2 ; 4H3S2  = H2S5  + 3H2S. 


COMPOUNDS  OF  SULPHUR  WITH  THE  HALOGENS. 

Sulphur  and  chlorine  unite  to  form  three  compounds:  SCI2,  SC1„  and 
S2CI2.  It  is  only  the  last  which  meets  with  any  practical  application. 

Sulphur  Dichloride — SC^ — is  produced  when  sulphur  monochloride,  S2CI2,  is  satu- 
rated with  chlorine  at  6°  to  10°  : 

S2CI2  + CI2  = 2SCI2. 

The  excess  of  chlorine  is  removed  by  conducting  a stream  of  dry  carbon  dioxide 
through  it. 

It  is  a dark-red  colored  lifpiid,  with  a specific  gravity  of  1.62;  boils  at  64°,  with 
j)artial  decomposition  into  S2CI2  and  ; the  dissociation  commences  at  ordinary  tem- 
peratures. 

Sulphur  Tetrachloride — SCI, — only  exists  at  temperatures  below  0°  C.  It  is 
formed  by  saturating  S2CI2  with  Cl  at  — 20°  to  — 22°.  Tlie  dissociation  commences 
at  — 20°,  and  is  comjdete  at  -[  6°.  It  yields  crystalline  compounds  with  some  chlo- 
ridc.s — e,  g.^  SnCl,,  AsCIj,  .SbCb. 


SULPHUR  MONOCHLORIDE. 


Ill 


The  most  stable  of  the  sulphur  chlorides  is 

Sulphur  Monochloride — S2CI2 — which  is  formed  when  chlorine  is 
conducted  over  molten  sulphur  contained  in  the  flask  C (Fig.  49).  It 
distils  over  and  condenses  in  the  receiver  D ; the  product  is  redistilled, 
to  obtain  it  pure.  [Fig.  49,  A:  chlorine  generator;  .•  wash-bottle  ; 
E : entrance  for  water  intended  to  chill  the  vapors.] 

Sulphur  monochloride  is  a reddish-yellow  liquid  with  a sharp  odor, 
provoking  tears,  having  a specific  gravity  of  1.68,  and  boiling  at  138°. 
Its  vapor  density  equals  135  (03  = 32),  corresponding  to  the  molecular 


formula  S2CI2.  It  fumes  strongly  in  the  air,  and  is  decomposed  by  water 
into  sulphur  dioxide,  sulphur  and  hydrochloric  acid  : 

2S2CI2  -f  2H2O  ==  SO2  -f  4HCI  + 3S. 


Sulphur  monochloride  dissolves  sulphur  readily  and  serves  in  the  vulcan- 
ization of  caoutchouc. 

Bromine  forms  analogous  compounds  with  sulphur.  S.^Brj  is  a red  liquid,  boiling  at 
190-200°.  Compounds  of  iodine  and  sulphur  are  not  definitely  known. 


II2 


INORGANIC  CHEMISTRY. 


1 


3.  SELENIUM. 

Atom  : Sc  = 79.1.  Molecule:  Scj  158.2  (at  1400°), 

Tliis  element  is  not  very  abundant  in  nalure,  and  is  only  found  in 
small  (luantities,  jirincijially  in  certain  iron  i>yritcs  (in  Sweden  and 
Jjohemia).  Upon  roasting  this  ore  of  iron,  for  the  prejiaration  of  sul- 
l)huric  acid,  selenium  settles  out  in  the  chimney  dust  or  in  the  deposit  of 
the  lead  chambers  (comiiare  Sulphuric  Acid),  and  was  found  there  by 
Berzelius  in  the  year  1817.  It  was  called  selenium  (frskrjurj,  moon)  to 
indicate  its  relation  to  the  already  known  tellurium  {Jcllus,  earth). 

Like  sul{)hur,  selenium  forms  different  allotropic  modifications.  Avior- 
phoiis  selenium,  obtained  by  the  reduction  of  selenium  dioxide  (SeOJ  by 
means  of  suli)hur  dioxide  (SOp,  is  a reddish-brown  powder,  soluble  in 
carl)on  bisulphide,  and  has  a specific  gravity  of  4. 26.  Selenium  crystallizes 
from  carbon  bisiiljihide  in  brownish-red  crystals  of  specific  gravity  4.5. 
The  solution  of  potassium  selenide  is  brown-red,  and  when  it  is  ex])osed  to 
the  air,  black  leaf-like  crystals  of  selenium  (with  si)ecific  gravity  4.8)  sepa- 
rate. These  are  isomorphous  with  suljdmr  and  insoluble  in  carbon  bisul- 
])hide.  Upon  suddenly  cooling  fused  selenium  it  solidifies  to  an  amorphous, 
glassy,  black  mass,  which  is  soluble  in  carbon  bisulphide  and  hasa  sjiecific 
gravity  of  4.28.  When  selenium  (amorphous)  is  heated  to  90-100°,  its 
temperature  suddenly  rises  above  200° ; it  is  converted  into  a crystal- 
line, dark-gray  mass  with  a specific  gravity  of  4.8.  It  possesses  metallic 
luster,  conducts  electricity,  and  is  insoluble  in  carbon  bisulphide.  The 
crystalline,  insoluble  modification  is  obtained  by  slowly  cooling  the  molten 
selenium. 

Selenium  melts  at  217°,  and  boils  at  about  660°,  passing  into  a dark- 
yellow  vapor.  The  vapor  density  diminishes  regularly  with  increasing 
temperature  (similar  to  sulphur),  and  becomes  constant  at  1400°.  It 
then  equals  158;  the  molecule  of  selenium  at  1400°  consists  of  two 
atoms  (2  X 79- 1 = 158-2). 

Selenium  resembles  sulphur  very  closely  in  its  chemical  behavior.  It 
burns  in  the  air  with  a reddish-blue  flame,  forming  selenium  dioxide,  and 
emits  a peculiar  odor  resembling  rotten  horse-radish.  It  dissolves  with  a 
green  color  in  concentrated  sulphuric  acid,  and  forms  selenious  and  sul- 
phurous acids : 

Se  -f-  2H2S0^  = SeOj  T 2SO2  + 2H2O. 

Hydrogen  Selenide — H2Se — produced,  like  hydrogen  sulphide, 
from  iron  selenide  and  hydrochloric  acid — is  a colorless,  disagreeably 
smelling  gas  with  poisonous  action.  In  the  air  the  aqueous  solution 
I ecomes  turbid  and  free  selenium  separates. 

Willi  chlorine  selenium  forms  SeCl,  and  Se2Cl2,  perfectly  analogous  in  general  to  the 
sulphur  cornjiounds  hut  differing  fnim  them  in  that  selenium  tetrachloride  is  a solid  which 
sublimes  and  does  not  begin  to  decompose  until  at  about  200°. 


TETXURIUM. 


II3 


4.  TELLURIUM. 

Atom:  Te  = i27.  Molecule:  Te2  = 254  (at  1700°).* 

Tellurium  is  of  rare  occurrence,  either  native  or  in  combination  with 
metals.  It  is  associated  with  gold  and  silver  in  sylvanite,  with  silver  and 
lead  in  altaite,  and  with  bismuth  in  tetradymite.  It  is  found  principally 
in  Transylvania,  Hungary,  California,  Virginia,  Bolivia,  Brazil,  and  in 
the  volcanic  formations  of  the  Liparian  Islands. 

The  tellurium  precipitated  by  sulphurous  acid  from  a solution  of  tel- 
lurous  acid  is  a black  ])ovvder  of  specific  gravity  5.9.  It  is  silver-white 
when  fused,  of  a perfect  metallic  luster,  and  conducts  electricity  and 
heat.  It  crystallizes  in  rhombohedra,  having  a specific  gravity  6.4.  It 
fuses  at  452°  and  boils  at  1390°.  When  heated  in  the  air  it  burns,  with 
a bluish-gray  flame,  to  tellurium  dioxide  (Te02). 

The  vapor  density  of  tellurium  at  1400-1700°  has  been  discovered  to 
be  about  254,  corresponding  to  the  molecular  formula  Te2. 

Hydrogen  Telluride — H2Te — formed  by  the  action  of  hydrochloric 
acid  upon  zinc  telluride,  is  a colorless,  very  poisonous  gas,  with  disagree- 
able odor. 

Two  chlorides — TeC]2  and  TeCfi,  and  two  bromides,  TeBr2  and  TeBr^ — have  been 
formed.  The  tetrachloride  boils  at  380°.  As  its  vapor  density  corresponds  to  the  formula 
TeCb  the  tellurium  is  quadrivalent. 


SUMMARY  OF  THE  ELEMENTS  OF  THE  OXYGEN  GROUP. 

The  elements  oxygen,  sulphur,  selenium  and  tellurium  form  a natural 
grouj)  of  chemically  similar  bodies.  The  similarity  of  the  last  three  is 
especially  marked,  while  oxygen,  ])ossessing  the  lowest  atomic  weight, 
stands  somewhat  apart.  Among  the  halogens,  fluorine  exhibits  a similar 
deportment;  it  departs  somewhat  from  its  analogues,  chlorine,  bromine 
and  iodine.  Like  the  halogens  the  elements  of  the  oxygen  group  present 
a gradation  in  their  properties  corresponding  to  their  atomic  weights  : 

os  Se  Te 
Atomic  weights,  16  32.06  79.1  127. 

With  the  increase  in  the  atomic  weight  there  occurs  a simultaneous 
condensation  of  substance,  the  volatility  diminishes,  while  the  specific 
gravity  and  the  points  of  fusion  and  boiling  increase,  as  may  be  seen  in 
the  following  table. 


*The  atomic  weight  of  tellurium  was  first  made  128  and  subsequently  125.  But  the 
later  development  of  the  periodic  system  made  it  more  probable  that  it  was  even  lower 
than  the  atomic  weight  of  iodine  (126.8),  which  view  has  since  been  confirmed  exjieri- 
mentally  by  Brauner.  More  recently  results  have  been  obtained  which  argue  for  the 
higher  atomic  weight,  d'e  --  127.  (See  Staudenmaier,  Zeit.  f.  anorg.  Chem,  10  (1895), 
189.) 

10 


in()R(;ani(:  cukmistry. 


114 


Oxy(;kn. 

Sui.i’inm. 

SlU.KNIUM. 

IT'.I.I.UKIUM, 

Specific  gravity,  . . 

1. 124  (at — 181°) 

1.95-2.07 

4. 2-4. 8 

r,.4 

Melting  point,  . . . 

114-5° 

217° 

452° 

Boiling  j)oint,  . . 

— 181° 

440° 

600° 

1 390° 

Gas  density,  .... 

32.00 

64. 12 

158.2 

254 

Oxygen  is  a difYicultly  c()ercil)le  gas,  while  the  others  are  solids  at 
ordinary  tem})eratures.  We  must,  however,  bear  in  mind  that  sulphur, 
selenium  and  tellurium  in  a free  state  are  probably  composed  of  larger 
complex  atomic  groiii)S  (see  j).  106). 

Further,  with  rising  atomic  weight  the  metalloidal  i)asses  into  a more 
metallic  character.  Tellurium  exhibits  the  physical  properties  of  a 
metal;  even  selenium  possesses  metallic  properties  in  its  crystalline 
modification.  In  chemical  deportment,  however,  the  metalloidal  char- 
acter shows  scarcely  any  alteration. 

All  four  elements  unite  directly,  at  elevated  temperatures,  with  hydrogen,  to  com- 
pounds the  composition  of  which  is  expressed  by  the  general  formula  MH.^.  At  still 
higher  temperatures  these  derivatives  are  again  resolved  into  their  elements.  The  simi- 
larity of  water  and  the  hydrides  of  sulphur,  selenium  and  tellurium  is  restricted  to  their 
formulas.  They  are  entirely  different  in  their  chemical  nature.  Water  is  absolutely 
necessary  to  life,  while  hydrogen  sulphide,  selenide  and  telluride  are  dangerous  poisons. 
It  must  also  occasion  surprise  that  oxygen,  the  only  gaseous  member  of  the  group,  forms 
a derivative  with  hydrogen  which  is  liquid  at  the  ordinary  temperature  and  a solid  at  0°, 
whereas  the  hydrides  of  the  elements  sulphur,  selenium  and  tellurium,  volatile  with  diffi- 
culty, are  also  gaseous  and  are,  comparatively  speaking,  condensed  with  more  difficulty. 
It  was  indeed  observed  with  the  halogens  that  the  hydride  of  fluorine,  the  most  difficult 
to  condense,  possessed  the  highest  boiling  point,  and  we  learned  that  the  explanation  for 
this  was  that,  in  accordance  with  its  vapor  density,  hydrogen  fluoride  did  not  have  the 
formula  HP",  but  H^F2.  Water  also  differs  very  much  the7-mo-che77ncally  from  the  hydrides 
of  the  other  elements  of  the  group.  It  is  a strongly  exothermic  body,  while  hydrogen 
sulphide,  on  the  other  hand,  formed  from  hydrogen  and  solid  sulphur  with  the  evolu- 
tion of  very  little  heat  (H2,S  = about  4 Cal.),  and  selenium  and  tellurium  hydrides  are 
indeed  endothermic  bodies. 

This  is  a gradation  similar  to  that  observed  in  the  halogen  hydrides  (p.  65). 
In  accord  with  this  we  find  that  oxygen  will  displace  these  elements  from  hydrogen 
sulphide,  hydrogen  selenide,  and  hydrogen  telluride  when  in  aqueous  solution  with  the 
formation  of  water.  At  higher  temperatures  and  with  an  excess  of  oxygen  the  dioxides 
(SO2,  Se02)  result.  The  halogens  decompose  them  more  readily  than  oxygen. 


NITROGEN  GROUP. 

This  group  consists  of  nitrogen,  ])hosphorus,  arsenic,  antimony,  and 
bismuth,  d'lie  last  ])ossesses  a decidedly  metallic  character.  It  does 
not,  like  the  other  four  elements,  form  gaseous  derivatives  with  hydrogen, 
d'he  gases  argon,  helium,  metargon,  kry])ton,  xenon,  and  neon,  occurring 
in  the  air,  will  l)e  described  in  connection  with  nitrogen. 


NITROGEN. 


1.  NITROGEN. 

Atom  : N = 14.04,  Molecule  : N2  = 28.08. 

Nitrogen  exists  free  in  the  air,  four-fifths  by  volume  are  nitrogen  and 
one-fifth  oxygen.  In  combination,  it  is  chiefly  found  in  the  ammonium 
and  nitric  acid  compounds,  as  well  as  in  many  organic  substances  of  the 
animal  and  vegetable  kingdoms. 

Until  1894  it  was  thought  possible  to  isolate  nitrogen  from  the  air 
by  depriving  the  latter  of  its  second  constituent — oxygen.  This  is 

effected  by  bodies  capable  of  ab- 
sorbing oxygen  without  acting 
upon  the  nitrogen,  e.  g , phos- 
phorus, hepar,  alkaline  solutions 
of  pyrogallol  and  heated  copper. 

The  experiment  can  be  most 
easily  and  simply  performed  in 
the  following  manner:  Several 
pieces  of  phosphorus  are  placed 
in  a dish  swimming  on  water, 
ignited,  and  a glass  bell-jar 
placed  over  them  (Fig.  50).  In 
a short  time,  when  all  the  oxy- 
gen is  absorbed  from  the  air,  the 
phosphorus  will  cease  burning ; 
the  phosphorus  pentoxide  pro- 
duced dissolves  in  the  water,  and  the  residual  gas  consists  of  almost  pure 
nitrogen  : its  volume  will  equal  four-fifths  of  the  air  taken.  Another 
procedure  consists  in  conducting  air  through  a red-hot  tube  filled  with 
copper  turnings;  the  copper  unites  with  the  oxygen  and  pure  nitrogen 
escapes.  The  portion  of  air  remaining  after  these  experiments  was  con- 
sidered to  be  nitrogen  until  1894  when  Lord  Rayleigh  and  W.  Ramsay 
proved  to  the  universal  surprise  that  so-called  “ atmospheric  nitrogen  ” 
was  a mixture  of  nitrogen  and  argon.  Lately,  other  very  probably 
elementary  gases  have  also  been  discovered  in  it. 

Pure  nitrogefi  can  only  be  obtained  from  nitrogenous  chemical  com- 
pounds. The  following  is  a simple  method  to  this  end:  Heat  ammonium 
nitrite  in  a small  glass  retort;  this  decomposes  the  salt  directly  into 
water  and  nitrogen : 

NH^N02  = N2  T 2H2O. 

In  place  of  ammonium  nitrite  a mixture  of  potassium  nitrite  fKNO.^)  and  ammonium 
chloride  (NIbCl)  may  be  u.sed  ; upon  warming,  the.se  salts  yield,  by  double  decomposi- 
tion, potassium  chloride  and  ammonium  nitrite  (KNO.^  + NH^Cl  = NH^NOg  + KCl), 
which  latter  decomposes  further.  As  potassium  nitrite  usually  contains  free  alkali,  some 
potassium  bichromate  is  added  to  neutralize  the  same.  Practically,  the  .solution  consists 
of  I part  of  potassium  nitrite,  i part  of  ammonium  chloride,  and  i part  of  potassium 
bichromate  in  5 parts  of  water,  and  is  then  boiled  ; to  free  the  liberated  nitrogen  from 
every  trace  of  oxygen  the  gas  is  conducted  over  ignited  copper. 

The  action  of  chlorine  upon  aqueous  ammonia  produces  pure  nitrogen. 
Ammonia  is  a compound  of  nitrogen  and  hydrogen.  While  the  chlorine 


Fig,  50. 


INORGANIC  CHFMISTRY. 


I l6 


conihinc's  vvitli  the  hydrogen  of  tlie  ammonia  to  liydroclilorie  acid,  nitro- 
gen is  lilierated.  d'he  hydrochloric  acid  unites  witli  the  excess  ol 
ammonia  to  ammonium  chloride  (NH^Clj.  'I'he  following  e<juations 
express  the  reactions : 


and 


2NII3  4-  3CI2  = Nj -f  GIICl 


GHCl  4 6NII3  = 6NH,C1. 

Animoiiiuin 

cliloridc. 


d'he  apiiaratus  jiictured  on  ]>.  50  will  serve  to  carryout  the  experiment. 
'The  disengaged  chlorine  is  conducted  through  a Woulff  wash-bottle  c(m- 
taining  ammonia  water,  the  free  nitrogen  being  collected  over  water. 

In  thi.s  expernnent  the  greate.st  care  should  he  exerci.se(l  tluit  an  excess  of  chlorinw  is 
not  conducted  into  the  solution,  because  its  action  ui)on  the  aminoniuin  chl(;ride  will 
cause  the  formation  of  an  exceedingly  explosive  body,  nitrogen  chloride,  NCI3  (p.  133;, 
which  separates  in  oily  drops. 

Pt'operiies  - — Nitrogen  is  a colorless,  odorless,  tasteless  gas.  One  liter 
of  it  weighs  I *2507  grams  at  0°  and  760  mm.  pressure.  It  is  therefore  0.97 
times  as  heavy  as  air.  Its  critical  temjierature  lies  near  — 146^^,  and  its 
critical  pressure  equals  35  almosjdieres  (p.  487.  Liquid  nitrogen  is  color- 
less, boils  under  a jtressure  of  one  atmosphere  at  — 194°,  at  — 225°  under 
a pressure  of  4 mm.,  and  has  a specific  gravity  of  0.885  — ^94°)- 

solidifies  at  — 214°. 

In  its  chemical  de])ortment  it  is  extremely  inert,  combining  directly 
with  but  few  elements,  e.  g.,  with  oxygen  and  hydrogen  under  the  in- 
fluence of  the  electric  S})ark ; with  magnesium  and  other  metals  at  more 
elevated  temperatures,  and  with  lithium  at  the  ordinary  temperature.  It 
does  not  support  combustion  or  respiration  ; a burning  candle  is  extin- 
guished and  animals  are  suffocated  by  it.  This  is  not  due  to  the  activity 
of  the  nitrogen,  but  to  the  absence  of  oxygen — a substance  which  cannot 
be  dispensed  with  in  combustion  and  respiration.  The  presence  of  nitro- 
gen in  the  air  moderates  the  strong  oxidizing  property  of  the  pure  oxygen. 


THE  ATMOSPHERE. 

The  air,  or  the  envelo])e  encircling  the  earth,  the  atmosphere  (ar/ioc — 
vapor;  frepaTpa — ball,  sphere),  consists  principally  of  a mixture  of  nitrogen 
and  oxygen;  it  always  contains,  in  addition,  slight  and  variable  quanti- 
ties of  aqueous  vapor,  carbon  dioxide,  and  traces  of  other  substances,  as 
accidental  constituents  (p.  123).  Recently  the  gases  argon,  helium,  met- 
argo7i,  neon,  krypton,  and  xenon  have  been  discovered  in  it.  They  occur 
in  small  amounts,  but  are  constant  constituents.  The  pressure  exerted 
by  the  air  is  measured  by  a column  of  mercury  which  holds  it  in  a 
state  of  ecpiilibrium  ; the  height  of  the  barometric  column  at  the  sea- 
level  and  0°  C.  e(]uals,  upon  an  average,  760  millimeters.  As  i c.c.  of 
mercury  weighs  13.6  grams,  76  c.c.  will  equal  1033.6  grams,  and  the  last 


THE  ATMOSPHERE. 


IT7 

number  would  indicate  the  pressure  which  the  column  of  air  exerts  upon 
one  square  centimeter  of  the  earth’s  surface. 

One  cubic  centimeter  of  air  weighs,  according  to  recent  experiments, 
under  normal  conditions,  0.00129276  gram  ; 1000  c.c.,  therefore,  or  one 
liter,  would  weigh  1.29276  grams.  As  one  liter  of  water  weighs  1000  grams, 
air  is  consequently  773  times  lighter  than  it.  Airis  14.4  times  heavier  than 
hydrogen.  Its  density  is  28.95  referred  to  O2  — 32.  Its  specific  volume 
is  773.4,  i.  e.,  I gram  of  air  at  0°  and  760  mm.  pressure  occupies  773.4 
c.c.  (pp.  45,  79). 

The  liquefactio7i  of  air — or  at  least  a portion  of  it — can  be  accomplished 
by  the  methods  described  on  p.  48.  Dewar  and  also  Olszewsky  chilled 
air,  by  means  of  boiling  ethylene,  almost  to  its  critical  temperature 
( — 140°),  and  then  liquefied  it  by  a simultaneous,  corresponding  increase 
of  pressure  (75  atmospheres).  Another  less  expensive  method,  hence  well 
adapted  for  technical  purposes,  utilizes  the  great  reduction  in  temperature 
sustained  by  strongly  compressed  gases  when  suddenly  released  Dom 
pressure.  In  the  apparatus  constructed  almost  simultaneously  by  Linde 
(Munich),  Tripler  (New  York)  and  Hampson  (London)  the  gas  under 
slight  pressure  and  cooled  by  sudden  expansion  is  directed  around  a cur- 
rent of  gas  passing  in  an  opposite  direction  under  high  pressure,  whereby 
the  temperature  is  eventually  lowered  below  the  critical  temperature. 

Liquid  air  is  colorless.  It  is  turbid,  owing  to  the  solid  carbon  dioxide 
from  which  it  can  be  freed  by  filtration.  Its  boiling  point  (according  to 
Dewar)  is  — 190°,  at  which  temperature  almost  pure  nitrogen  is  evolved. 
The  specific  gravity  of  freshly  prepared  liquid  air  is  0.9951  referred  to 
water  at  4°.  It  contains  much  more  oxygen  than  the  gas.  Recently 
prepared  it  holds  as  much  as  54  percent,  of  oxygen  by  weight,  while  the 
gas  form  has  only  23  per  cent,  of  oxygen  by  weight.  Hence  the  name  liquid 
air  is  not  an  entirely  correct  designation.  The  oxygen-content  of  liquid 
air  increases  by  preservation.  This  may  be  accomplished  by  means  of 
Dewar  bulbs  or  double-walled  tin  or  wooden  boxes  the  air  space  of  which 
is  filled  out  with  silk,  etc.  It  will  finally  contain  as  much  as  94  per  cent, 
by  weight  of  oxygen.  This  gradual  accumulation  of  the  gas  in  the  liquid 
air  is  attributable  to  the  fact  that  under  the  ordinary  pressure  nitrogen 
boils  at  — 194°;  oxygen,  however,  at  — 184°.  There  is  here  the  greatest 
technical  possibility,  that  by  boiling  out  nitrogen  from  liquid  air,  almost 
pure  oxygen  can  be  prepared  on  a large  scale. 

The  effect  of  low  temperatures  upon  the  physical  properties  of  bodies 
and  upon  the  course  of  chemical  reactions  can  be  well  shown  by  means 
of  liquid  air.  Carbon  dioxide  and  acetylene  solidify  in  it.  The  solid 
acetylene  may  be  ignited;  it  then  burns  away  like  paraffin.  Liquid  air 
immediately  solidifies  mercury  and  renders  it  malleable.  Alcohol  at 
once  forms  drops  in  liquid  air  and  soon  becomes  hard  and  crystalline. 
The  hand  may  be  plunged  for  a short  period  into  the  liquid  of  — 190° 
without  experiencing  any  ill  effects,  because  it  is  at  once  surrounded  by 
a protecting  film.  Dewar  claims  that  the  color  of  many  bodies  is  changed 
if  they  are  immersed  in  liquid  air.  Red  mercuric  oxide,  iodide  and 
sulphide  (cinnabar)  appear  yellow  in  color,  while  the  yellow-green  nitrate 
of  uranium  appears  white.  Mention  has  already  been  made  that  chemical 


ii8 


INORGANIC  CHEMISTRY. 


transpositions  do  not  occur,  or  at  least  very  slowly,  at  the  temperature  of 
lifpiid  air  (p.  28). 

History. — It  is  well  known  that  the  air  was  formerly  regarded  as  an  element — a 
simple  substance.  However,  observations  were  made  very  early  which  argue  for  the 
very  opposite.  As  early  as  the  ninth  century  certain  chemists  knew  that  the  metals,  when 
heated  in  air — i.e.,  by  calcination — increased  in  weight,  and  some  of  them  correctly 
attributed  this  to  the  absor[)tion  of  certain  air  particles  l)y  the  hot  metal.  Because  niter, 
like  air,  accelerates  combustion,  1 looke  (seventeenth  century)  suspected  that  both  contained 
an  ingredient  .serving  for  combustion.  It  was  especially  the  English  chemists  of  the  .seven- 
teenth century  who  busied  themselves  with  researches  in  this  direction.  Mayow  in  particular 
deserves  mention.  lie  contended  that  the  glowing  metal  united  with  the  “ spiritus  nitro- 
aerus”  of  the  air.  He  demonstrated  that  by  respiration  as  well  as  by  combustion  the 
volume  of  air  standing  over  water  was  diminished,  and  that  the  residuum — the 
destroyed  air — was  no  longer  available  for  respiration  or  combustion.  However,  the.se 
germs  of  the  correet  idea  of  combustion  phenomena  could  not  develop  and  grow  under 
the  predominance  of  the  peculiar  views  then  extant.  Becher,  a German  chemist,  began 
about  1700  to  teach  that  the  phenomenon  of  combustion  originated  in  a peculiar  volatile 
and  escaping  earth  or  kind  of  sulphur.  His  pupil  Stahl,  about  1720,  developed  this 
thought  into  a theory  or  doctrine,  which  held  almost  exelusive  sway  until  near  the  closing 
third  of  the  century.  Stahl  called  the  substance,  which  escaped  in  combustion, 
(combustible).  He  failed  to  note  that  which  became  the  main  subject  of  investigation 
with  the  English  chemists — that  the  burnt  metal  had  increased  in  weight,  and  he 
explained  that  all  combustible  bodies  consisted  of  phlogiston  and  a non-combu.stible  sub- 
stance. Thus,  sulphur  consisted  of  phlogi.ston  and  sulphuric  acid,  iron  of  ferric  oxide  and 
phlogiston,  etc.,  etc.  In  the  process  of  combustion  the  phlogiston  escapes  as  flame  and 
the  non-combustible  portion  remains.  The  reduction  of  a metallic  oxide  by  another 
metal  or  by  a combustible  body  {e.g.,  carbon)  depends,  therefore,  upon  the  passage  of 
phlogiston  from  l^ie  reducing  to  the  reduced  body  ; hence  the  terms — phlogisticated  and 
dephlogisticated.  The  first  corresponds  to  reduction  and  the  second  to  oxidation.  In  this 
manner,  because  of  the  absence  of  more  accurate  experiments,  the  phenomena  of  combustion 
could  be  explained  with  some  probability  and  apparent  certainty.  And  Stahl  became  the 
founder  of  a chemical  theory  which  lasted  for  more  than  a half  century,  and  in  its 
decadence  found  its  most  ardent  advocates  among  the  best  and  most  celebrated  English 
chemists  of  that  period.  Bayen  (1774)  showed  that  mercuric  oxide  was  reduced  without 
the  addition  of  phlogiston  (i.  e.^  by  heating  without  the  addition  of  carbon),  and  this  led 
Lavoisier  to  his  experiments  upon  the  absorption  of  air  in  the  calcination  of  the  metals.  He 
melted  tin  in  a large,  closed,  air-tight  flask.  He  had  previously  determined  its  weight 
as  well  as  that  of  the  tin.  When  the  latter  had  become  coated  with  a thick  layer  of  oxide 
the  apparatus  was  allowed  to  cool  and  was  again  weighed.  Its  weight  was  the  same  as 
at  first,  but  when  the  flask  was  opened,  air  rushed  in  and  the  weight  increased.  Hence 
the  tin,  by  calcination,  had  abstracted  something  from  the  air  and  had  not,  as  required 
by  the  phlogiston  theory,  given  anything  to  it.  About  this  same  period  Scheele  and 
Priestley  discovered  oxygen  (p.  80),  Rutherford  (1742)  had  again  described  destroyed 
air,  i.  e.,  nitrogen,  and  Scheele,  by  a series  of  excellent  experiments,  demonstrated  the 
composition  of  air  and  the  difference  between  nitrogen  (aer  mephiticus,  aer  vitalis)  and 
carbon  dioxide  (aer  aereum).  But  he  believed  in  the  phlogiston  theory.  It  was 
Lavoisier,  during  the  .same  period,  who  clearly  indicated  and  showed  experimentally  the 
role  of  oxygen  in  combustions  and  oxidations,  and,  shortly  after,  proved  the  elementary 
nature  of  nitrogen.  Lavoisier  named  nitrogen  azote  (from  life,  and  a,  privative), 
from  which  we  get  the  .symbol  Az  u.sed  in  France  for  nitrogen.  Chaptal  was  the  first  to 
give  nitrogen  the  name  nitrogenium  (whence  the  symbol  N),  because  it  is  a constituent 
of  niter  (nitrum).  How  Engli.sh  chemists  have  again  extended  our  knowledge  of  the 
constituents  of  the  air  to  a most  unexpected  degree  may  be  gathered  from  pp.  I15-123. 
Coirn)are  Berzelius’  Lehrbuch  der  Chemie,  5 Auil.  (1843),  I,  140. 

Quantitative  Cottiposition  of  Air. — Its  composition  is  expressed  by  the 
quantity  of  oxygen,  argon  and  nitrogen  contained  in  it,  as  its  remaining 
admixtures  are  more  or  less  accidental  and  variable. 


THE  ATMOSPHERE. 


II9 

Boussingault  and  Dumas  (181 1)  determined  the  accurate  weight  com- 
position of  the  air  (nitrogen  and  argon  were  of  course  calculated  as  ‘‘at- 
mospheric nitrogen  ”)  by  the  following  experiment : A large  balloon,  V, 
with  a capacity  of  about  20  liters  (Fig.  51),  is  connected  with  a porcelain 
tube,  a,  b,  filled  with  metallic  copper.  Balloon  and  tubes,  closed  by 
stop-cocks,  are  previously  emjitied  and  weighed  apart.  The  bent  tubes, 
A,  B,  and  C,  contain  caustic  potash  and  sulphuric  acid,  and  serve  to  free 
the  air  undergoing  analysis  from  aqueous  vapor,  carbon  dioxide,  and  other 
impurities.  The  porcelain  tube,  filled  with  copper,  is  heated  to  a red 
heat,  and  by  carefully  opening  the  stop-cocks  ?/,  r,  and  r a slow  current 
of  air  is  allowed  to  enter  the  empty  balloon  V.  The  impurities  are  given 
up  in  the  bent  tubes,  and  all  the  oxygen  absorbed  by  the  ignited  copper, 
forming  cupric  oxide,  so  that  only  pure  nitrogen  enters  V.  Now  close  the 


cocks  and  weigh  the  balloon  and  porcelain  tube,  a,  b.  The  increase  in 
weight  of  the  latter  represents  the  quantity  of  oxygen  in  the  air;  the 
increase  in  V the  quantity  of  nitrogen.  In  this  manner  Dumas  and 
Boussingault  found  that  in  100  parts  by  weight  of  air  there  are  con- 
tained : 


Nitrogen, 7^-995  parts  by  weight. 

Oxygen, 23.CX35  “ “ “ 

Air, 100,000  “ “ “ 


Asa  liter  of  oxygen  weighs  1.4291  grams  and  a liter  of  atmospheric 
nitrogen  1.25 71  grams  the  volume  composition  of  air  would  be: 


Oxygen, 20.8  parts  by  volume. 

Nitrogen, 79.2  “ “ “ 

Air, loo.o  “ “ “ 


The  density  of  air  is  28.95  (P-  45)  referred  to  oxygen  etpial  to  32. 


I 20 


INORGANIC  CHEMISTRY. 


olf^ 


Fig.  52. 


the  original 


'Fhc  voliiinc  coiii])osition  of  air  may  1)C  directly  fcjiind  l)y  means  of  the 
al)sori)tiometer.  'I'he  latter  is  a tube  carefully  graduated,  and  sealed  at 
one  end.  This  is  filled  with  mercury,  and  air  allowed  to  enter;  the 
volume  of  the  latter  is  determined  by  reading  off 
^ the  divisions  on  the  tube.  Now  introduce  into 

jijl  the  tube,  through  the  mercury,  a i)latinum  wire 

i | having  a ball  of  i)hos])horus  attached  to  the  end 

(Fig.  52),  or  a ball  of  coke  saturated  with  an 
alkaline  solution  of  jiyrogallic  acid.  The  phos- 
phorus absorbs  the  oxygen  of  the  air,  and  only 
nitrogen  remains,  the  volume  of  which  is  read  off 
by  the  graduation. 

The  eudiometric  affords  greater  accuracy. 

It  is  dependent  ii])on  the  combustion  of  the  oxy- 
gen with  hydrogen  in  a eudiometer.  Air  and 
hydrogen  are  introduced  into  the  eudiometer,  and 
the  electric  sjiark  then  ])assed  through  the  wires. 

All  the  oxygen  in  the  air  combines  with  a ])ortion 
of  the  hydrogen  to  form  water.  On  cooling,  the 
aqueous  vapor  condenses  and  a contraction  in 
volume  occurs.  Assuming  that  we  had  taken  100 
volumes  of  air  and  50  volumes  of  hydrogen,  and 
that  the  residual  volume  of  gas,  after  allowing  for 
all  corrections  (p.  121),  equaled  87.15;  then  of 
150  volumes  of  mixed  gas,  62.85  volumes  disap- 
peared in  the  formation  of  water.  As  the  latter  results  from  the 
union  of  i volume  of  oxygen  and  2 volumes  of  hydrogen,  the 
TOO  volumes  of  air  employed  in  the  analysis  therefore  contained 
6 2^8^  = 20.95  volumes  of  oxygen.  Hence  air  consists  (accord- 
ing to  the  determination  of  Regnault  and  Bunsen)  of 

79.05  volumes  Nitrogen. 

20.95  “ Oxygen. 

100.00  “ Air. 

A eudiometer  {evbui,  fair  weather,  and  iiirpov,  measure — for- 
merly it  was  thought  that  there  was  some  connection  between 
the  quantity  of  oxygen  in  the  air  and  the  weather)  is  an  absorp-  piQ.  53. 
tiometer,  with  two  })latinum  wires  fused  into  its  closed  end. 

'The  i)assage  of  the  electric  spark  from  wire  to  wire  causes  the  explosion 

53)-* 

Numerous  analyses  show  that  the  composition  of  the  air  everywhere  on 
the  earth’s  surface  is  constant.  The  most  recent  and  exhaustive 
researches  of  Kreusler,  Hempel,  Morley,  Rayleigh,  and  Leduc  indicate 
it  to  be  generally  as  follows: 


* For  furtlier  .sliidy  of  gasometric  methods  consult  Rob.  Bunsen,  (lasometrische  Meth- 
oden,  2 Aull.,  1877;  Wallher  irem])el,  Clasanalytische  Melhoden,  2 Aufk,  1890; 
Clemens  Winkler,  Lehrhuch  der  Ic^clinischen  (lasanalyse,  2 Aull.,  1892. 


THE  ATMOSPHERE. 


I 21 


Nitrogen,  78.06  parts  by  volume  ; 
Oxygen,  21.00  “ “ “ 

Argon,  0.94  “ “ “ 


75.5  parts  by  weight. 
23.2  “ “ “ 

1.3  “ “ “ 


100.00 


100.0  “ “ 


Measuring  Gases. — The  volume  of  gases  is  influenced  by  pressure,  temperature,  the 
moisture  contained  in  them,  and  to  a slight  degree  by  their  chemical  nature.  To  com- 
pare statements  of  gas  volumes,  they  must  be  recalculated  to  normal  conditions,  i.  e.,  it 
must  be  indicated  what  volume  is  occupied  by  the  gas  at  0°  and  760  mm.  barometric 
pressure.  The  pressure  exerted  by  the  mercury  column  of  a barometer  at  a definite  tem- 
perature depends  not  only  on  its  height  but  also  upon  the  intensity  of  gravity,  which  in 
turn  varies  with  the  latitude  and  sea-level.  Hence  it  has  been  agreed  to  refer  the  “nor- 
mal conditions”  to  45°  geographical  latitude  and  the  sea-level.  The  normal  volume 
Vo  (at  760  mm.  and  0°)  is  calculated,  according  to  Boyle-Gay-Lussac’s  laws,  from  the 
volume  V,  which  a gas  occupies  at  pressure  p and  t°  by  the  equation  : 


pv  = po  Vo  (l  + at) 

to  be 

V . p 

760  (l  -f  at)‘ 

This  answers  only  for  dry  gases.  Any  quantity  of  gas  occupies  less  space  when  it  is  dry 
than  when  it  is  moist,  because  the  tension  of  the  aqueous  vapor  counteracts  the  atmos- 
pheric pressure. 

The  moisture  may  be  removed  by  introducing  into  the  gas  a ball  of  coke  saturated 
with  sulphuric  acid,  which  dries  it.  It  is  more  convenient,  however,  to  make  the  cor- 
rection of  the  gas  volume  in  the  following  manner  : Water  is  brought  in  contact  with  the 
gas  to  be  measured,  in  order  to  perfectly  saturate  it  with  aqueous  vapor  ; the  gas  is  then 
measured  and  its  normal  volume  calculated  by  the  above  formula,  after  deducting  from 
the  observed  pres.sure  p the  number  of  millimeters  corresponding  to  the  tension  of  the 
aqueous  vapor  for  the  given  temperature  (p.  90).  If  the  aqueous  tension  at  t®  be  repre- 
sented by  s (mm.  mercury)  we  finally  reach  the  equation  : 

V ==  V • (P  — s) 

° 760  (i  + 0.003665  . t)‘ 

Another  form  may  be  given  the  expression  of  the  two  laws  relating  to  gases  : 
pv  = Po  Vo  (i  -f  Ot), 


if  the  temperature  be  counted  not  from  the  melting  of  ice  forward  but  from  the  absolute 
zero  point  of  temperature,  which  can  be  developed  from  the  following  considerations. 
According  to  Gay-Lussac  the  volume  of  a gas  increases,  for  every  degree  of  rise  in  tem- 


perature, the  0.00367  or  ^ of  the  volume  it  occupied  at  0°.  If  the  temperature  be 

lowered  from  0°  downward,  and  the  law  of  Gay-Lussac  continues  to  hold  force,  then  the 
volume  of  the  gas  at  — 273°  equals  zero.  This  degree  of  temperature,  which  has  been 
approached  to  within  30°,  answers  for  the  zero  point  of  ab.solute  temperature.  It  is  indi- 
cated by  the  letter  T.  The  relation  between  the  degrees  of  ab.solute  temperature  and 
those  of  ordinary  centigrade  degrees  is  expres.sed  by  t = T — 273.  If  this  be  introduced 
into  the  preceding  erjuation  we  obtain 


pv  ^ 


Pn  Vo 

273 


T, 


or,  if  be  made  equal  to  R,  then  we  have 
’273  i ’ 


pv  = R . T. 


The  value  of  R does  not  depend  upon  (he  chemical  composition  of  the  gas,  but  .solely 
upon  the  units  of  measure  chosen  ff)r  p and  v.  d'be  law  of  Avogadro  can  also  be  given 
cxpres.sion  through  tliis  ef|uation  if,  following  Horstmann’s  suggestion,  consideration  be 
given  to  the  volumes  which  molecular  quantities  of  the  gases  occupy, — /.  e.  the  volume 

II 


22 


INORGANIC  CHEMISTRY. 


occu})ic(l  by  one  molecule  or  2.02  f^rams  of  hydrogen,  32  grams  of  oxygen,  70.9  grams 
of  chlorine,  36.46  grams  of  hydrogen  chloride.  'I'his  volume  is,  Cf)n.se(iuently,  22.4 
liters  (p.  98).  In  this  case  R is  the  .same  for  all  gases.  If  the  volume  be  measured  in 
cubic  centimeters,  and  the  j)res.sure  in  grams  j)cr  .sfiuare  centimeter,  then,  as  v = 224(X), 
p would  ecpial  1033.6  (p.  116),  T - 273,  and  R 84800. 

'I'he  entire  expre.ssion  for  the  three  laws  would  then  read 

pv  = 84800  T. 

We  shall  see  that  this  equation  also  posse.sses  great  value  for  solutions. 

The  I^oyle-Ciay-Lussac  law  is  not  an  ab.solutely  correct  expre.ssion  for  the  behavior  of 
gases.  At  very  great  pre.ssures  they  can  be  less  compres.sed  than  wovdd  correspond  to  the 
letpiirements  of  the  law.  At  low  pre.ssures,  on  the  other  hand,  gases,  hydrogen  excepted, 
can  be  more  strongly  compressed  than  the  law  re(|uire.s.  'J'he.se  variations  justify  the 
theory  propounded  by  the  Hollander,  van  der  Waals  ( 1873).  As  the  density  of  the  gas 
increases  the  attraction  between  its  molecules  becomes  greater  and  the  outward  ))ressure 
grows  less;  indeed,  the  diminution  is  inversely  proj)nrtional  to  the  .scjuare  of  the  volume 
of  the  gas-mass.  Accordingly,  the  observed  luessure,  when  com])ared  with  that  demanded 

l)y  the  law,  is  reduced  about  the  value  , in  which  a is  a constant  corresponding  to  the 

attraction  between  the  molecules — the  cohesion  of  the  gases.  On  the  other  hand,  the 
space  remaining,  with  the  greater  density,  for  the  motions  of  the  molecules  is  le.ss  than 
the  observed  volume,  becau.se  we  must  deduct  from  the  latter  the  .s{)ace  actually  occupied 
by  the  molecules  themselves.  In  calculation  the  ob.served  pressure  must  therefore  be 
increased  and  the  ob.served  volume  must  be  diminished.  This  is  indicated  in  the  equa- 
tion of  van  der  Waals,  mentioned  on  p.  49  : 

(P  4-  “0  (v-b)  = RT. 
or 

(P  + ■ vV)  ('■  - b)  = (I  -b  a)  (I  - b)  ,* 

which  not  only  answers  for  gases  but  also  for  liquids.  The  critical  data  of  a gas  may 
also  be  calculated  from  it. 

From  the  great  constancy  of  its  composition  air  was  supposed  to  be  a 
chemical  compound,  consisting  of  nitrogen  and  oxygen.  This  supposi- 
tion is,  however,  opposed  by  the  following  circumstances:  All  chemical 
compounds  contain  their  constituents  in  atomic  quantities,  which  is  not 
the  case  with  air.  In  the  mixing  of  nitrogen  and  oxygen  to  form  air 
there  is  neither  disengagement  nor  absorption  of  heat,  which  is  always 
observed  in  chemical  compounds.  Further,  the  air  absorbed  by  water  or 
other  solvents  jiossesses  a composition  different  from  the  atmospheric; 
this  is  due  to  the  unequal  solubilities  of  nitrogen  and  oxygen  in  water. 
'The  air  expelled  from  water  iqion  application  of  heat  consists  of  34.9 
volumes  of  oxygen  and  65.1  volumes  of  nitrogen  (Bunsen).  These  facts 
indicate  that  air  is  not  a chemical  compound,  but  a mechanical  mixture 
of  its  two  constituents  (see  Liquid  Air). 

I'he  great  con.staucy  iu  the  comj^osition  of  our  atmo.sjdicre  is  due  chiefly  to  the  fact 
that  there  is  a constantly  renewed  mixture  produced  by  the  unceasing  air  currents,  by 
wind.s  and  storms,  rising  of  the  warmer  layers  and  the  sinking  of  those  which  have 
become  cooler.  The  mutual  diffusion  of  gases  comes,  therefore,  into  consideration.  The 


-x-As  K 


" ' , a.Hl  v=  I. 

273  ‘ 


GASES  RECENTLY  DISCOVERED  IN  THE  ATMOSPHERE. 


123 


gas  molecules  possess,  as  is  now  generally  acknowledged,  a direct,  progressive,  energetic 
movement,  and  can  therefore  diffuse  without  limitation  into  space  if  by  contact  with  other 
molecules  they  are  not  deflected  from  or  arrested  in  their  course.  This  is  the  reason  that 
two  gases  (or  liquids)  in  immediate  contact  mix  generally  with  one  another.  Another 
kind  of  diffusion need  not  be  considered  so  far  as  concerns  the  constant 
composition  of  the  atmosphere,  but  may  be  given  as  a conclusion  to  what  has  already 
been  said.  If  two  gases  are  separated  by  a permeable  diaphragm  or  membrane,  that  one 
possessing  the  lower  density  will  traverse  the  septum  the  more  rapidly. 

The  following  experiment  very  clearly  illustrates  this  : In  the  open  end  of  an  unglazed 
clay  cylinder  (as  used  in  galvanic  elements)  there  is  fixed  a glass  tube  about  one  meter 
long,  its  open  end  terminating  in  a dish  containing  water  ; the  cylinder  and  tube  are  filled 
with  air.  Over  the  porous  cylinder  is  placed  a wider  vessel  filled  with  hydrogen.  The 
latter  presses  faster  into  the  cylinder  than  the  air  escapes  from  it ; the  air  in  the  tube  and 
cylinder  is  displaced  and  rises  in  the  water  in  bubbles.  \\'hen  the  separation  of  gas 
ceases,  tube  and  cylinder  are  almost  filled  with  pure  hydrogen.  On  removing  the  larger 
hydrogen  vessel  the  gas  will  escape  much  more  rapidly  into  the  external  air  than  the 
latter  can  enter  the  cylinder  ; the  internal  pressure  will  therefore  be  less  than  the  external, 
and  water  ascends  in  the  glass  tube. 

In  addition  to  nitrogen  and  oxygen,  air  constantly  contains  aqueous 
vapor  and  carbon  dioxide  (CO2)  in  very  small  quantities.  The  presence 
of  the  former  can  readily  be  recognized  by  the  fact  that  cold  bodies  are 
covered  with  dew  in  moist  air.  Its  quantity  depends  on  the  temiterature 
and  corresponds  to  the  vapor  tension  of  water  (see  p.  90).  One  cubic  meter 
of  air  perfectly  saturated  with  aqueous  vapor  contains  22.5  grams  of 
water  at  25°  C.  ; on  cooling  to  0°  17.1  grams  separate  as  rain.  Gen- 
erally the  air  contains  only  50-70  per  cent,  of  the  quantity  of  vapor 
necessary  for  complete  saturation.  The  amount  of  moisture  in  it  is  either 
determined  according  to  physical  methods  (hygrometer),  or  directly  by 
weighing.  To  this  end  a definite  quantity  of  air  is  conducted  through  a 
tube  filled  with  calcium  chloride  or  sulphuric  acid,  and  its  increase  in 
weight  determined. 

To  detect  the  carbon  dioxide  in  the  air,  conduct  a portion  of  the 
latter  through  solutions  of  barium  or  calcium  hydroxides,  and  a turbidity 
will  ensue.  To  determine  its  quantity,  pass  a definite  and  previously 
dried  amount  of  air  through  a weighed  potassium  hydrate  tube,  and  ascer- 
tain the  increase  in  weight  of  the  latter.  Ten  thousand  parts  by  volume 
of  atmospheric  air  contain,  ordinarily,  from  3.0. to  4 parts  by  volume  of 
carbon  dioxide.  (See  Ber.  30  (1897),  1450.) 

Besides  the  four  ingredients  just  mentioned,  air  usually  contains  small 
quantities  of  ozone,  hydrogen  peroxide,  ammonium  salts  (ammonium 
nitrite),  the  newly  discovered  gases  argon  and  heliiu?i,  and  probably 
hydrogen.  Finally,  air  contains  microscopic  germs  of  lower  organisms ; 
they  are  generally  found  in  the  lower  air  strata,  and  their  presence  influ- 
ences the  processes  of  decay  and  fermentation  of  organic  substances. 


GASES  RECENTLY  DISCOVERED  IN  THE  ATMOSPHERE. 

Lord  Rayleigh  in  1892,  while  engaged  in  an  exhaustive  research  upon  the  density  of 
the  elementary  gases,  incidentally  observed  that  a liter  of  nitrogen  isolated  from  the  air 
weighed  1.2571  grams  under  normal  conditions,  whereas  a liter  of  the  same  gas  pre- 
pared from  ammonia  or  nitric  acid  weighed  1.2507  grams.  The  determined  purpose  of 
ascertaining  the  reason  for  this  difference  in  the  third  decimal  led  Lord  Rayleigh  and 


124 


INORGANIC  CHEMISTRY. 


W,  Ramsay  in  1894  to  the  brilliant  and  snr])risinfi;  discovery  of  argon — a new  conslitnont 
of  tlie  air,  the  composition  of  wliich,  after  the  masterly  investigations  of  emiiK-nt  chemists, 
ai)peared  to  conceal  from  us  no  furtlier  enigmas.  Vet,  in  Marclj  of  1895,  as  W.  Ramsay 
sought  ft)r  additional  .sources  of  argon,  lie  found  heliiim,  a second  ajiparently  elementary 
gas,  which,  through  a jiortion  of  its  spectrum,  had  been  known  as  a constituent  of  the  sun. 
Prior  to  this  it  had  not  been  certaiidy  known  to  be  pre.se nt  on  the  earth.  It  is  a very 
subordinate  constituent  of  the  air,  where  it  was  first  ob.served  by  Kayser.  Ramsay 
obtained  it  by  heating  minerals  containing  uranium. 

The  two  gases  to  which  reference  has  just  been  made  differ  from  all  other  gases  known 
to  us  by  the  total  absence  of  chemical  activity.  'Phus  far  no  other  substance  has  been 
made  to  react  with  them.  Hence  the  name  Argon,  from  iiv  ipyov,  without  action.  'They 
al.so  differ  from  other  gases  in  their  real  atomic  structure  ; their  atoms  have  not  com- 
bined to  molecules— a condition  which  heretofore  has  been  ob.served  oidy  with  other 
substances  at  very  high  temperatures  (ji.  79).  'Phis  is  evident  from  the  relation  of  the 
specific  heat  at  constant  jiressure  to  tliat  at  constant  volume,  which  in  the  case  of  mon- 
atomic gases  eijuals  1.67,  whereas  with  polyatomic  gase.s — tho.se  built  uj)  molecularly — 
the  value  lies  between  i and  5.  Compare  ().  IT  Meyer,  'Phe  Kinetic  Theory  of  Ga.ses 
(2d  ed.,  1895-1899). 

Recently,  Ramsay  availed  himself  of  the  remarkably  developed  refrigatory  apparatus 
and  found  four  apparently  elementary  monatomic  ga.ses  in  air,  while  seeking  for  the  one 
between  argon  and  helium.  He  permitted  large  quantities  of  licjuid  argon  and  liquid  air 
to  boil  away  gradually,  and  then  examined  the  more  volatile  portions  and  those  portions 
not  so  readily  volatile  and  more  difficultly  condensed.  Thus  he  di.scoverecl  in  the  residue, 
by  the  evaporation  of  750  c.c.  of  liquid  air  to  Jo  c.c.,  the  gas  krypton  {Kpv-iTTor,  con- 
cealed), which  volatilizes  with  difficulty,  has  a density  or  atomic  weight  of  45  (O2  — 32), 
and  is  characterized  in  its  spectrum  by  a brilliant  red,  a yellow,  and  a green  line,  the 
last  lying  close  to  that  of  the  northern  light ; similarly  metargon  and  xenon  [^evoq, 
foreign),  both  of  which  are  present  in  the  heavier  portions  of  the  air  ; and  lastly  the  non- 
condensible, light  neon  [veo^,  new),  rich  in  lines,  with  the  density  20,  and  probably  the 
sought-for  gas  lying  between  argon  and  helium.  No  one  of  these  four  ga.ses  has  been 
obtained  in  the  pure  state  ; therefore  only  argon  and  helium  will  be  more  fully  discussed. 
(See  Ramsay,  Ber.  31  (1898),  31  ii.) 


ARGON. 

Atom  : A = 40. 

The  air  contains  about  0.935  cent,  by  volume  of  argon,  while  “atmospheric  nitro- 
gen ” contains  1.183  per  cent,  by  volume.  We  inhale  daily  about  20  liters  of  this  enig- 
matical gas  which  has  remained  concealed  for  so  long  a time.  It  is  pre-sent  in  many 
mineral  waters  and  gases  from  springs,  e.g.,  those  of  Bath,  Cauteret,  Voslau  and  Wild- 
bad.  It  is  liberated  in  small  quantities,  however,  together  with  helium,  upon  heating 
certain  minerals  (especially  those  containing  uranium) — cleveite,  broggerite,  uraninite, 
and  has  been  obtained  in  the  same  manner  from  one  meteorite.  Frequently  the  gases 
occluded  in  rock-salt  contain  argon. 

'Pwo  methods  may  be  pursued  in  separating  argon  from  the  air  : 

1.  Air — freed  from  oxygen,  aqueous  vapor  and  other  admixtures — “atmospheric 
nitrogen” — (p.  119)  is  conducted  over  red-hot  magnesium  filings,  which  ab.sorb  the  nitro- 
gen, forming  magnesium  nitride  : N2  -j-  Mg.^  :=  Mg3N2  (p.  ii6),  while  the  argon  is  scarcely 
acted  upon.  The  nitrogen  can  be  more  quickly  absorbed  by  lithium  or  by  a heated  mix- 
ture of  magnesium  and  calcium  oxide  or  finely  divided  calcium  : 

Mg  CaO  = MgO  -f  Ca. 

Ram.say  employed  this  method. 

2.  Lord  Rayleigh  mixed  “atmospheric  nitrogen,”  the  mixture  of  nitrogen  and  argon, 
with  oxygen  and  allowed  the  induction  spark  to  pass  continuously  through  this  mixture 
in  the  presence  of  caustic  ])ota.sh.  Alkaline  nitrite  is  formed  in  this  way  from  the  nitro- 
gen, oxygen  and  alkali  (sec  Nitrous  Acid,  ])]X  205,  207),  while  the  residue  of  argon  and 
oxygen  is  conducted  over  ignited  co]q)er  to  remove  the  latter.  Cavendi.sh,  as  early  as 
1785,  observed  that  there  remained  .some  “ atmo.s])heric  nitrogen  ” which  was  not  con- 
verted into  alkaline  nitrile  ; he  did  not,  however,  pursue  this  observation  further. 


COMPOUNDS  OF  NITROGEN  WITH  HYDROGEN. 


125 


The  argon  tlius  prepared  contains  helium  and  neon  as  the  lighter  and  inetargon, 
krypton  and  xenon  as  tlie  heavier  admixtures.  It  is  then  liquefied  by  means  of  licpiid 
air  and  purified  by  fractional  distillation.  Its  density  referred  to  C)^  32  is  39.914  or  in 

round  numbers  40  ; it  is  identical  with  its  atomic  weight.  Under  normal  conditions  a 
liter  of  argon  weighs  1.780  grams.  At  very  low  temperatures  it  congeals  to  an  ice-like 
mass,  which  melts  at  — 189.5°  ^o  a colorless  liquid.  It  boils  at  — 185°  ; its  critical 
temperature  is — 121°  and  its  critical  pressure  50.6  atmospheres.  Argon  is  approxi- 
mately times  as  soluble  in  water  as  nitrogen  : 100  volumes  of  water  dissolve  4 

volumes  of  argon  at  12°.  Hence  argon  accumulates  in  the  gases  of  rain-water.  The 
spectrum  distinguishes  argon  very  certainly  from  nitrogen  and  other  substances  [com- 
pare : Z.  f.  phys.  Ch.  1895,  16,  344;  Z.  f.  anorg.  Ch.  18  (1898),  222  ; Chem.  Central- 
blatt  70  (1899),  I,  469;  also  Mugdan  : Argon  and  Helium,  two  new  gaseous  ele- 
ments, Stuttgart,  1896]. 

HELIUM. 

Atom  : He  = 4. 

Helium  is  as  inactive  chemically  as  argon.  While  the  latter  is  widely  distributed  on 
the  earth,  helium  is  one  of  the  rarest  of  terrestial  substances.  It  occurs,  however,  with 
hydrogen  in  immense  masses  in  the  photosphere  of  the  sun  and  other  brilliant  fixed 
stars.  Norman  Lockyer  detected  it  as  early  as  1868  by  means  of  the  spectroscope  in  the 
chromosphere  of  the  sun,  and  in  1869  he  and  Frankland  named  it  Helium.  In  1892 
Palmieri  observed  its  terrestial  occurrence  for  the  first  time  while  studying  spectroscopically 
a substance  which  had  been  thrown  out  by  Vesuvius.  But  it  was  first  in  March,  1895, 
that  W.  Ramsay  found  helium  while  he  was  engaged  in  seeking  for  sources  of  argon. 
Hitherto  only  a portion  of  the  helium  spectrum  (line  Dg)  had  been  known.  The  gas 
helium  (together  with  hydrogen,  carbon  dioxide,  nitrogen  and  probably  also  argon)  is 
evolved  when  certain  rare  minerals,  usually  consisting  of  salts  of  uranium,  yttrium  and 
thorium  (^.  g.,  cleveite,  uraninite,  broggerite,  monazite)  are  heated  alone  with  dilute  sul- 
phuric acid  (i  : 8)  or  with  sodium  bisulphate.  It  was  similarly  obtained  from  a meteorite. 
Kayser  found  it,  in  small  quantities  it  is  true,  in  the  air  and  the  later  investigations  of 
Ramsay  and  Travers  have  confirmed  this  observation.  Helium,  like  argon,  occurs  in  the 
gases  from  springs,  e.  g.,  in  those  from  Wildbad,  in  that  of  Adano  near  Padua,  as  well 
as  in  the  gaseous  exhalations  of  Tuscany. 

Helium  is  monatomic  (see  above).  Its  atomic  weight  (its  density)  equals  about  4. 
Dewar  has  shown  that  it  can  be  liquefied  by  means  of  liquid,  vaporizing  hydrogen,  so 
that  at  present  all  gases,  with  the  exception  of  neon,  can  be  liquefied.  It  is  less  soluble 
in  water  than  any  other  gas  : loo  volumes  of  water  at  18°  take  up  but  0.73  volume  of 
helium.  It  is  as  inactive  as  argon.  Travers  asserts  that  when  a powerful  electric  dis- 
charge is  sent  through  a Pliicker  tube  (see  Spectrum  Analysis)  filled  with  helium  the  latter 
is  absorbed  by  the  platinum  electrodes  and  is  again  liberated  from  the  same  on  the  appli- 
cation of  heat.  This  recalls  the  behavior  of  hydrogen  and  palladium  ; it  may  be  useful 
in  separating  helium  and  argon.  Five  brilliant  lines  are  prominent  in  the  spectrum  of 
helium  : one  each  in  the  red,  the  yellow,  the  green,  the  blue  and  the  violet.  The  yellow 
line  Dg  lies  close  to  the  two  sodium  lines,  Dj  and  D^,  toward  the  violet  end  of  the 
.spectrum. 


COMPOUNDS  OF  NITROGEN  WITH  HYDROGEN. 

Ammonia  is  the  most  important  compound  of  nitrogen  and  hydrogen. 
It  has  l^een  known  for  tlie  longest  time.  Since  1889  Th.  Curtins  has 
added  four  otlier  derivatives  of  these  two  elements  to  it : hydrazine,  N^H^, 
and  hydrazoic  acid,  N.Jl,  as  well  as  the  ammonium  and  hydrazine  salts 
of  the  latter:  NJI,  (=  NHg  + NgM)and  NHI,  (=N.^H,  + NgH).  See 
Per.  29  ( 1896),  759.  Oxvaminonia  or  hydroxylaminc,  NHgO,  discovered 
by  W.  J.ossen  in  1865,  will  be  discussed  after  ammonia. 


26 


INORGANIC  CHEMISTRY. 


1.  AMMONIA. 

Mulcculc  : NUa  :^  17.07. 

Ammonia  occurs  in  tlie  air  in  combination  with  some  acids,  in  natural 
waters  and  in  the  eartli,  but  always  in  small  (juantities,  Ihiestley  first 
studied  it  carefully  and  called  it  alkaline  air.  berthollet  determined  its 
composition  in  17CS5.  d'lie  formation  of  ammonia  by  the  direct  union 
of  nitrogen  and  hydrogen  occurs  under  the  inlluence  of  thesilent  electric 
discharge.  Its  comiiounds — ammonium  salts — are  freciuently  ])roduced 
under  the  most  varying  conditions.  Thus  ammonium  nitrate  is  formed 
by  the  action  of  the  electric  spark  iijion  moist  air: 

K,  + O + 2ll,0  NII^NOg. 

Aminoiiium 

nitrate. 

Ammonium  nitrite,  NH^NOj,  is  said  to  be  formed  in  every  combus- 
tion in  the  air;  and  in  the  electrolysis  of  water.  Further,  ammonium 
salts  are  ])roduced  in  the  solution  of  many  metals  in  nitric  acid,  in  con- 
setpience  of  a reduction  of  the  acid  by  the  liberated  hydrogen  : 

HNO3  + 81 1 311,0  + NIIj. 

The  following  conversion  of  nitrogen  into  ammonia  deserves  considera- 
tion : Nitrogen  unites  with  magnesium  at  a red  heat  to  magnesium  nitride, 
MggN,,  and  the  latter  is  energetically  decomposed  by  water  with  the 
evolution  of  ammonia: 

MggN,  + 311,0  = 3MgO  + 2NII3  (p.  124). 

Nitric  acid  in  the  form  of  salts,  in  alkaline  solution,  is  reduced  by 
nascent  hydrogen  to  ammonia  (see  Nitric  Acid). 

Ammonia  is  produced  in  large  quantities  in  the  decomposition  and 
dry  distillation  of  nitrogenous  organic  substances.  Even  as  late  as  the 
last  century  the  bulk  of  the  ammonium  chloride  (the  most  important  salt 
technically),  was  obtained  from  camel’s  dung  or  decayed  urine.  Its  orig- 
inal name  Sa/  armoniacimi — Armenian  salt — was  confounded  later  with 
the  designation  for  Egyptian  rock-salt — Sal  ammoniacum.  In  the  prepa- 
ration of  illuminating  gas  by  the  distillation  of  coal,  ammonia  appears  as 
a by-i)roduct  and  may  be  obtained  by  combining  it  with  sulphuric  or 
hydrochloric  acid.  This  method  is  used  almost  exclusively  at  present  for 
its  production. 

To  ])re[)are  ammonia  heat  a mixture  of  ammonium  chloride  and  slaked 
lime  in  a glass  or  iron  flask  : 

2NIFCI  f Ca(On),  ==  CaCl,  + 2ll,0  + 2NH3. 

Aininoiiium  Calcium 
chloride.  hydroxide. 

d'he  disengaged  ammonia  gas  is  collected  over  mercury,  as  it  is  readily 
soluble  in  water  (see  ]>.  58,  Fig.  33).  For  ])erfect  drying  conduct  it 
through  a vessel  filled  with  burnt  lime  (CaO).  ('alcium  chloride  is  not 
applicable  for  this  ])urpose,  as  it  combines  with  the  gas.  In  consequence 
of  its  levity,  ammonia,  like  hydrogen,  may  be  collected  by  disi)lacing  the 
air  in  inverted  ves.sels. 


AMMONIA. 


127 


Physical  Propc7'ties. — Ammonia  is  a colorless  gas  with  a suffocating, 
characteristic  odor.  Its  density  is  0.59  (air  = i).  Under  a pressure  of 
6.5  atmospheres  (at  10°  C.),  or  by  cooling  to  — 40°  C.,  under  ordinary 
pressure  it  condenses  to  a colorless  mobile  liquid  with  a specific  gravity 
of  0.623  at  0°,  solidifies  at  — 85°,  and  melts  again  at  — 75°.  Liquid 
ammonia  has  recently  been  introduced  into  commerce. 

Ammonia  gas  may  be  condensed,  just  like  chlorine.  Take  ammonium  silver  chloride 
(AgCl . 2NH3),  obtained  by  conducting  ammonia  over  silver  chloride,  and  enclose  it  in 
a tube  with  a knee-shaped  bend  (p.  51,  Fig.  32).  'I'he  limb  containing  the  compound  is 
now  heated  in  a water-bath,  while  the  other  limb  is  cooled.  The  compound  is  decom 
posed  into  silver  chloride  and  ammonia,  which  condenses  in  the  cooled  limb  (h'aradayj. 

Ammonia  gas  dissolves  very  readily  in  water,  with  the  liberation  of 
heat.  One  part  of  water  at  0°  and  760  mm.  pressure  absorbs  1146  volumes 
(=  0.875  parts  by  weight);  at  20°,  739  volumes  (=  0.526  parts  by 
weight)  of  ammonia.  At  16°  and  760  mm.  pressure  100  parts  by  volume 
of  water  absorb  60  parts  by  weight  of  ammonia.  The  specific  gravity  of 
a 34.95  per  cent,  solution  at  15°  is  0.882;  henc  e a liter  of  it  contains 
308.3  grams  of  ammonia.  When  a long  glass  tube,  closed  at  one  end 
and  filled  with  ammonia,  has  its  open  end  placed  in  water,  the  latter 
rushes  up  into  the  tube  as  it  would  into  a vacuum ; a ])iece  of  ice  melts 
rapidly  in  the  gas.  The  aqueous  solution  possesses  all  the  properties  of 
the  free  gas,  and  is  called  Liquor  ammonii  caustici.  d'he  greater  the 
ammonia  content  the  less  will  be  the  specific  gravity  of  the  solution.  All 
the  gas  escapes  on  the  application  of  heat. 

When  the  condensed  liquid  ammonia  evaporates  it  absorbs  a great  amount  of  heat, 
and  answers,  therefore,  for  the  production  arti- 
ficially of  cold  and  ice  in  Carre’s  apparatus. 

The  simplest  form  of  the  latter  is  represented  in 
P'ig.  54.  The  iron  cylinder  A is  filled  about 
half  with  a concentrated  aqueous  ammonia  solu- 
tion, and  is  connected,  by  means  of  the  tubes 
from  b,  with  the  conical  vessel  Aj  in  the  middle 
of  which  is  the  empty  cylindrical  space  E.  The 
entire  internal  s])ace  of  A and  A is  hermetically 
shut  off.  A is  heated  upon  a charcoal  fire  until 
the  thermometer  a,  in  it,  indicates  130°  C., 
while  A' is  cooled  with  water.  In  this  way  the 
gaseous  ammonia  is  expelled  from  the  aqueous  ^ 
solution  in  A^  passes  through  b,  in  which  most 
of  the  water  runs  back,  and  condenses  to  a 
liquid  in  B,  of  the  receiver  F.  The  cylinder  A 
is  removed  from  the  fire,  cooled  with  water  and 
the  vessel  A>,  constructed  of  thin  .sheet-metal 
and  filled  with  water,  placed  in  the  cavity  A,  FiG.  54. 

which  is  surrounded  with  a poor  conductor, 

e.  g.,  felt.  The  ammonia  condensed  in  B evaporates,  and  is  reab.sorbed  by  the  water  in 
A.  Ily  this  evaporation  a large  fjuantity  of  heat,  withdrawn  from  A’ and  its  surroundings, 
becomes  latent ; the  water  in  1)  freezes. 

'I'he  method  of  Carre  for  the  artificial  production  of  ice  has  acquired  great  apjdication 
in  the  arts;  recently,  however,  ice  machines  have  been  introduced,  'i'hese  are  driven 
by  liquid  ammonia  (Lindel,  or  by  licjuid  sulphur  dioxide  and  carbon  dioxide  (Pictet) 
(compare  pp.  48,  117).  The  method  of  Windhausen,  depending  upon  the  expansion 
of  compressed  air,  is  much  used. 


128 


INORGANIC  CIIEMISIRY. 


Chemical  Properties. — A red  heat  or  the  continued  action  of  the  elec- 
tric spark  decomposes  ammonia  into  nitrogen  and  liydrogen.  On  con- 
ducting ammonia  gas  over  heated  sodium  or  potassium,  tlie  nitrogen 
combines  with  these  metals  and  hydrogen  escapes  : 

NII3  -f- 3K  NK3 -p  3li. 


Magnesium  unites,  when  heated  in  an  atmosphere  of  ammonia,  with 
the  nitrogen  of  the  latter,  giving  a bright  light: 

3Mg  -I-  2NII3  _ Mg3N3  }-  3II3. 

Ammonia  will  not  burn  in  the  air;  in  oxygen,  however,  it  burns  with 
a yellow  flame  : 

2NIl3  + 3().=  N3  + 3ll30; 

ammonium  nitrite  and  nitrogen  dioxide  are  formed  simultaneously. 
When  a mixture  of  ammonia  and  oxygen  is  ignited  it  burns  with  explo- 
sion. 

To  show  the  combustion  of  ammonia  in  oxygen,  proceed  as  follows:  A 
glass  tube,  through  which  ammonia  is  conducted,  is  brought  with  oxygen 
into  a vessel,  bringing  the  opening  of  the  latter  near  a flame  at  the  moment 
of  the  introduction  of  the  glass  tube.  In  contact  with  oxygen,  the  am- 
monia gas  ignites  and  continues  to  burn  in  it. 

The  following  experiment  (of  Kraut)  shows  the  combustion  of  ammonia 
very  conveniently.  Place  a somewhat  concentrated  ammonia  solution  in 
a beaker  glass;  heat  over  a lamp,  until  there  is  an  abundant  disen- 
gagement of  gas,  and  then  run  in  oxygen,  by  means  of  a tube  dipped 
into  the  liquid.  Upon  approaching  the  mixture  with  a flame,  it  ignites 
with  a slight  explosion.  The  ignition  may  be  induced  without  a flame, 
by  sinking  a glowing  platinum  spiral  into  the  mixture ; we  then  have  a 
number  of  slight  explosions.  The  glass  is  filled  at  the  same  time  with 
white  vapors  of  ammonium  nitrite  (NH^N02)  ; later,  when  oxygen  pre- 
dominates, red  vapors  of  nitrogen  dioxide  (NO^)  and  nitrous  acid 
appear. 

If  chlorine  gas  be  conducted  into  the  vessel  with  ammonia,  it  immedi- 
ately ignites  and  continues  to  burn  in  the  latter,  with  the  production  of 
white  fumes  of  ammonium  chloride  (NH^Cl).  The  chlorine  combines 
with  the  hydrogen  of  the  ammonia,  with  separation  of  nitrogen,  and 
yields  hydrochloric  acid,  which  unites,  with  the  excess  of  ammonia,  to 
form  ammonium  chloride  : 

NII3  -f  3CI  = 3riCl  + N 

and 

3NII3  ~p  3IICI  3NII,C1. 

Chlorine  reacts  similarly  u]3on  acpieous  ammonia  (p.  ti6). 

In  gaseous  form,  as  well  as  in  solution,  ammonia  possesses  strong  basic 
])roj)erties ; it  bines  red  litmus  jiaper  and  neutralizes  acids,  forming  salt- 
like comj)onnds  with  them,  which  are  very  similar  to  the  salts  of  the 
alkalies — sodium  and  potassium.  'The  following  illustrates  the  similarity. 


QUANTITATIVE  COMPOSITION  OF  AMMONIA. 


I 29 


NII3 

+ 

IICl 

= NII^Cl 

Ammonium 

chloride. 

KCl. 

Potassium 

chloride. 

2NH3 

+ 

H2SO4 

= (Nil,)., SO, 

Ammonium 

sulphate. 

K3SO,. 

Potassium 

sulphate. 

NHj 

+ 

H3S 

= NII,SH 

Ammonium 

sulphydrate. 

KSH. 

Potassium 

sulphydrate. 

In  these  ammonia  derivatives  NH^  plays  the  role  of  a metal.  Hence, 
to  show  its  similarity  to  sodium,  potassium  and  other  metals,  the  group 
(NH^)  has  been  designated  ammonium  and  its  compounds,  ammonium 
salts.  The  latter,  when  acted  on  by  strong  bases,  yield  ammonia  gas : 

2NH,C1  + CaO  = 2NH3  + CaCb  + 

The  metallic  character  of  the  ammonium  group  is  also  confirmed  by  the 
existence  of  the  ammonium  amalgam.  Therefore,  the  ammonium  deriva- 
tives will  be  considered  with  the  metals. 

Thermo-chemical  Deportjnent. — The  heat  of  formation  of  ammonia  from 
hydrogen  and  nitrogen  equals  12  Cal.  When  ammonia  gas  is  dissolved 
in  much  water  8.4  Cal.  are  set  free,  so  that  the  heat  of  formation  of 
ammonia  from  its  elements  in  dilute  aqueous  solution  equals  20.4  Cal.  : 

(N,H3  - gas)  = 12.  (NH3,Aq)  = 8.4.  (N,H3,Aq)  = 20.4. 

The  great  heat  of  solution  of  gaseous  ammonia  explains  why  ice  will  melt  in  the 
same  {p.  127). 

The  explosibility  of  a mixture  of  ammonia  and  oxygen  is  accounted  for  by  the  follow- 
ing great  heat  disengagement : 

2NH3  + 30  = 3H3O  + N,  . . (+  147.6  Cal.) 

(24  Cal.)  (3  X 57-2  Cal.) 

The  action  of  chlorine  upon  gaseous  or  aqueous  ammonia  is  also  very  energetic  : 

NH3  gas  + 3CI  = 3HCI  gas  + N . . . (+  54  Cal.) 

(12  Cal.)  (66  Cal.) 

NIl3-dissolved  + 3CI  = 3TICl-dis.solved  + N . . . (+  97.5  Cal.) 

(20.4  Cal.)  (3X39-3  Cal.) 

When  there  is  an  excess  of  ammonia  the  hydrochloric  acid  combines  with  it  to  form 
ammonium  chloride  (NII3  -f-  HCl  = NII^Cl),  and  the  heat  disengagement  is  thereby 
further  increased. 


QUANTITATIVE  COMPOSITION  OF  AMMONIA.  ATOMIC  WEIGHT  OF 

NITROGEN. 

The  quantitative  analysis  of  ammonia  shows  that  it  consists  of  i.oi 
parts  of  hydrogen  and  4.68  parts  of  nitrogen;  hence  we  conclude  that 
the  atomic  weight  of  nitrogen  is  4.68  or  a multiple  of  it  (see  p.  69)  : 


II  = I 
N = 4.68 


2II  = 2.02 

N = 9.34 

NII2  - 11.36 


3II  = 3.03 

N = 14.04 


Nil 


5.68 


Nil, 


17.07 


T30 


INORGANIC  CHKMISTRV. 


As  tlic  density  of  ammonia  ccjiials  i 7.07  (O.^ — 32),  its  molecular  weight 
would  e(}ual  17.07.  In  17.07  parts  of  ammonia  there  are  3.03  parts,  and, 
therefore,  three  atoms  of  hydrogen.  'That  the  14.04  partsol  nitrogen  united 
with  tliem  corres])ond  to  one  atom  of  nitrogen  is  a consc-(|uence,  as  never 
less  than  1^.04  parts  of  that  element  are  present  in  the  niotecnlar  weije;ht  of 
any  nitrogen  derivative.  The  density  of  nitrogen  ecjuals  28.08  ( Oj  = 32 ) ; 
therefore,  the  molecule  of  nitrogen  consists  of  two  atoms  (N^).  'This  is 
also  concluded  from  the  volume  ratios,  as  we  shall  soon  see,  occurring  in 
the  decomposition  of  ammonia  (comi)are  p.  76). 

From  the  molecular  formulas  NH.,  and  N2  it  follows,  further,  that  i 
volume  of  nitrogen  and  3 volumes  of  hydrogen  form  2 volumes  of  am- 
monia gas,  or  that  2 volumes  of  ammonia  decompose  into  3 volumes  of 
hydrogen  and  i volume  of  nitrogen,  corresponding  to  the  molecular 
equation  : 

N2  + 3H2  = 2NII3. 

I vol.  3 vols.  2 vols. 

The  following  exiieriments  jirove  these  conclusions: 

1.  Decompose  an  acpieous  ammonia  solution,  mixed  with  salt  (NaCl) 
to  increase  its  power  of  conductivity,  in  a Hofmann’s  apparatus  (j).  77), 
by  the  galvanic  current.  Hydrogen  will  sei)arate  at  the  negative 
and  nitrogen  at  the  positive  ])ole;  the  former  will  have  three  times  the 
volume  of  the  latter  as  soon  as  the  solution  is  saturated  with  gases. 

2.  Pass  electric  (induction)  sparks  through  dry  ammonia  gas  contained 
in  a eudiometer,  or  the  apj^aratus  represented  in  Fig.  47  (p.  98).  In  this 
way  the  ammonia  is  decomposed  into  nitrogen  and  hydrogen,  the  volume 
of  which  is  twice  as  large  as  that  of  the  ammonia  employed.  That  3 
volumes  of  hydrogen  are  present  in  the  mixture  for  every  volume  of  nitro- 
gen is  easily  shown  by  the  eudiometric  method,  by  burning  the  hydrogen 
with  oxygen  (p.  99). 


2.  HYDROXYLAMINE  (OXYAMMONIA). 

NH3O  = NHoOH. 

This  compound  was  discovered  (by  Lossen  in  1865),  in  the  reduction 
of  ethyl  nitrate  by  tin  and  hydrochloric  acid,  in  the  form  of  its  salts  and 
in  aqueous  solution.  Lobry  de  Bruyn  first  obtained  it  anhydrous  and 
in  a solid  form  in  1891.  It  is  produced,  too,  by  the  action  of  tin  upon 
dilute  nitric  acid,  and  by  tin  and  hydrochloric  acid  upon  all  the  oxygen 
compounds  of  nitrogen.  In  all  these  reactions  it  is  the  hydrogen  elimi- 
nated by  the  tin  which,  in  statu  7iaseendi,  reduces  the  nitric  acid : 

IINO3  -f  311.2  =3  NH3O  -f  2H2O. 

To  prepare  hydroxylamine  treat  ethyl  nitrate  f 120  grams)  with  granulated  tin  (400 
grains)  and  hydrocddoric  acid  (800-1000  c.c.  of  specific  gravity  1.19,  mixed  with  three  times 
its  volume  of  water).  'I'he  metal  should  he  completely  di.s.solved.  The  .solution  diluted 
to  twice  its  volume  is  treated  with  hydrogen  .suli)hide  to  precipitate  the  tin.  The  filtrate 
from  llie  tin  sul|)hide  is  evaporated  and  the  hydroxylamine  hydrochloride,  NII3O  . IICl, 
extracted  from  the  residue  with  hot  alcohol. 

I lydroxylamiiie  hydrochloride  is  most  easily  formed  by  the  interaction  of  hydrochloric 
acid  and  fulminating  mercury  (see  Organic  Chemistry). 


HYDROXYl. AMINE. 


I3I 

Ilydroxyhmiine  suli)liate  is  technically  prepared  by  healing  potassium  hydroxylamine- 
disulphonate  with  water  to  100-130°,  when  hydroxylainine  sulphate  and  potassium  sul- 
phate are  produced  [see  Raschig,  Ann.  Chem.  241  (1887),  161]  : 

NII0(S03K)2  H 2II2O  = NIT3O.  II^SO^  -f  K^SCV 

Potassium  hydroxylamine-disulphonate  is  formed  by  the  interaction  of  potassium  nitrite 
and  acid  potassium  sulphite  in  the  cold  : 

KNO.2  + 2KHSO3  = NH0(S03K)2  -f  KOH. 

To  obtain  solid  hydroxylainine  decompose  the  hydrochloride  in  methyl  alcohol  solution 
by  a corresponding  quantity  of  sodium  methylate  ; hlter  from  the  sodium  chloride  which 
separates  and  distil  the  filtrate  under  greatly  reduced  pressure  [Briihl,  Ber.  26  (1893), 
III,  2508;  also  Lobry  de  Bruyn,  ibid.  27  (1894),  i,  967]. 

Hydroxylamine  crystallizes  in  colorless,  odorless  needles.  It  melts  at 
33°  and  under  22  mm.  pressure  boils  at  58°.  Its  specific  gravity  equals 
1.235.  rather  stable  at  temperatures  below  15°,  but  above  that  it 

rapidly  breaks  down  into  nitrogen,  nitrous  oxide,  nitrous  acid  and 
ammonia.  At  about  130°  and  often  at  much  lower  temperatures  the 
decomposition  is  accompanied  with  explosion.  It  absorbs  moisture 
rapidly  on  exposure  to  the  air. 

Hydroxylamine  is  very  similar  to  ammonia,  and  like  it  unites  directi}- 
with  the  acids  to  form  salts  : 

NH3O  + HCl  = NH3O.  HCl  = NH3(0H)C1. 

Hydroxylamine  hydrochloride  in  distinction  to  ammonium  chloride  is 
soluble  in  ale 'hoi.  It  passes  into  ammonium  chloride  when  allowed  to 
stand  exposed  to  the  air. 

On  adding  to  the  aqueous  solution  of  the  sulphate  of  hydroxylamine 
sufficient  barium  hydroxide  to  remove  all  the  sulphuric  acid,  an  aqueous 
solution  of  hydroxylamine  is  obtained,  which,  like  the  ammonia  solution, 
possesses  strong  basic  properties,  and  colors  red  litmus-paper  blue.  The 
solution  is,  however,  very  unstable,  and  readily  decomposes  into  water, 
ammonia,  and  nitrogen  : 

3NH3O  NH3  + 3H2O  + N2. 

Upon  the  application  of  heat  a portion  of  the  hydroxylamine  will  be 
carried  over  undecomposed  along  with  the  steam,  but  most  of  it  is 
decomposed.  The  hydroxylamine  solution  manifests  a strong  reducing 
action  ; it  preci])itates  metallic  silver  from  silver  nitrate,  white  mercurous 
chloride,  Hg2Cl2,  from  mercuric  chloride,  HgCl2,  and  cuprous  oxide  from 
cupric  salts. 

Owing  to  its  great  similarity  to  ammonia  and  its  various  reactions,  it  is 
supposed  that  hydroxylamine  represents  ammonia  in  which  an  hydrogen 
atom  is  rejilaced  by  the  hydroxyl  grouji  OH;  it  is  therefore  a compound 
of  the  latter  group  with  the  amido  grou]) : 

NII3O  = NII2.  OH. 

3.  Diamide,  or  Hydrazine,  N2TI4  ^ H2N  . Nir2,  a compound  of  two  amido-groups 
(NII2),  was  until  1889  only  known  in  its  numerous  organic  derivatives.  Since  then  Cur- 
tius  and  Ids  co-workers  have  exhaustively  investigated  a large  number  of  its  inorganic 
compounds,  and  in  1895  Bobry  de  Bruyn  succeeded  in  getting  the  free  diamide  or  hydra- 


132 


INORGANIC  CHEMISTRY. 


ziiie  unknown  until  then  [Her,  28  (1895),  30^5 J-  salts  were  first  ma«le  hy  an  indirect 

method  from  certain  orjranic  ‘ nitrogen  derivatives:  hy  the  decomposition  of  dia/.o-  or 
triazo-acetic  acid  uj)on  digesting  them  with  water  or  mineral  acids  [Curtins  and  Jay,  J.  f. 
prakt.  Ch.  [2]  39  (1889),  27J  ; by  heating  amidoguanidine  with  sodium  hy<lr(ixi<le 
[ Thiele,  Ann.  Chem.  270  (1892),  31],  etc.  .See  also  the  last  edition  of  Richter’s 
Organic  Chemistry.  Very  recently  diamide  has  been  made  fnmi  inorganic  compcnmds  ; 
by  the  reduction  of  nitric  oxide-potassium  sulphite  with  nascent  hydrogen  : 

K2SO3.  N2O2  -f  3II.2  = N.2II4.  ll.p  -f  K^SO^ 

[Duden,  Ber.  27  (1894),  3498]. 

Free  hydrazine,  like  free  hydroxylamine,  is  prepared  by  decomposing  its  hydrocliloride 
with  sodium  methylate  : 

N2II,.  IICl  + CIIgONa  = NaCl  + CII3OII  -f  N^H,. 


It  may  also  be  obtained  by  distilling  its  hydrate  under  reduced  pressure  with  barium 
oxide.  It  is  a colorless  liquid,  fuming  strongly  in  the  air,  boiling  at  113.5°  .solidify- 
ing at  0°.  It  melts  at  1.4°.  Its  specific  gravity  is  about  1.003  23°.  It  is  not  explo- 
sive and  is  stable  even  at  300°,  Its  hydrate,  II./J,  is  a strong  base,  very  much 

like  ammonia.  It  is  ])roduced  when  the  hydrazine  salts  are  decompo.sed  by  powerful 
bases.  It  is  characterized  by  great  reducing  power.  It  will  be  discus.sed  later  along 
with  the  diammonium  salts  in  concluding  the  alkaline  earth  metals. 

4.  Hydrazoic  Acid,  or  Azoimide,  N3II  = |1  /Nil,  was  also  di.scovered  by  Cur- 


tins in  1890  [Ber.  23  (1890),  3023].  It  was,  like  the  preceding  body,  first  isolated 
from  organic  derivatives,  and  indeed  from  tho.se  of  diamide  [especially  benzoyl  azoimide 
which  upon  digestion  with  sodium  alcoholate  yields  sodium  azoimide  and  ethyl  benzoate  : 


CeH5CON3  + C^IipNa  NaN3  + C.II.COOC^IIs, 


and  from  diazohippuramide  which  is  resolved  by  ammonia  into  ammonium  azoimide, 
hippuramide  and  ammonia  : 

C6H5CO-nh-ch2-co-nh-n=n-oh  -f  2NH3  = 

N3.  NH^  -f  H2O  -h  C6li5CO-NH-CH2-CO-Nil2. 

[A  description  of  these  investigations  can  be  found  in  the  Jahrbuch  der  Chemie  I (1891), 
91,  227  ; II  (1892),  229  ; also  Thiele,  Ann.  Chem.  270  (1892),  53,  and  Richter’s  Organic 
Chemistry.] 

The  following  are  real  inorganic  methods  for  the  preparation  of  hydrazoic  acid  or  its 
salts. 

Curtius  prepares  a dilute  aqueous  solution  of  the  acid  by  chilling  with  ice  the  red 
vapors  arising  from  heating  arsenious  oxide  and  nitric  acid  (see  Nitrous  Acid)  and 
adding  the  resulting  blue  liquid  gradually,  as  long  as  gas  is  not  evolved,  to  a cold  solution' 
of  hydrazine.  This  is  analogous  to  the  formation  of  silver  azoimide  (Angeli)  from  silver 
nitrite  and  hydrazine  sulphate  : 

N==N 

H2-N-NH2  -b  NO2H  1=  H2-N-N=N-0H  + H2O  = \ / -f  2H2O. 

Nil 


Wislicenus  as.serts  that  .sodium  azoimide  is  readily  obtained  by  heating  sodamide  to  150- 
240°  in  a current  of  nitrous  oxide  : 


NIl2Na  -f  N.2O  = NaNg  + II.2O. 


Ililute  aiiueous  .solutions  of  hydrazoic  acid  may  be  obtained  by  di.stilling  the  metallic 
azoimidcs  with  dilute  sul])lmric  acid.  By  further  distillation  of  the  solution  and  by 
means  of  calcium  chloride  the  acid  may  be  obtained  anhydrous  when  it  is  a colorless, 


NITROGEN  CHLORIDE. 


133 


mobile  liquid  with  an  unbearable  odor.  It  boils  at  37°,  volatilizes  with  steam, 
and  is  characterized  by  its  exceptionally  violent  explosibility.  In  this  respect  it  differs 
from  the  haloid  acids,  especially  hydrochloric  acid,  with  which  it  possesses  many  points 
of  similarity.  This  behavior  would  indicate  that  hydrazoic  acid  is  a strongly  endothermic 
body  : 

(N3,H)  dissolved  = —62  Cal. 

Its  alkali  salts  are  not  explosive.  Some  of  the  alkaline  earth  .salts  explode  when  heated, 
and  some  of  the  heavy  metals  w’hen  touched  or  struck.  These  salts  will  be  further 
treated  under  the  respective  metals.  [See  Curtius  and  Rissom,  Jr.  prakt.  Ch.  58  (1898;, 
261  ; also  Z.  f.  anorg.  Ch.  17  (1898),  18.] 


COMPOUNDS  OF  NITROGEN  WITH  THE  HALOGENS. 

NITROGEN  CHLORIDE. 

NCI3. 

As  we  have  seen,  nitrogen  is  liberated  when  chlorine  acts  upon  an 
excess  of  ammonia  (p.  116) ; when,  however,  the  chlorine  is  in  excess,  it 
acts  upon  the  previously  formed  ammonium  chloride  to  produce  nitrogen 
chloride : 

NH,C1  -f  3CI2  = NCI3  + 4HCI. 

For  the  preparation  of  a small  quantity  of  nitrogen  chloride,  dip  a 
flask  filled  with  chlorine,  open  end  down,  into  an  aqueous  ammonium 
chloride  solution  warmed  to  30°.  The  chlorine  is  absorbed,  and  heavy 
oil  drops  separate,  which  are  best  collected  in  a small  leaden  dish.  They 
contain  hydrogen  in  addition  to  nitrogen  and  chlorine.  They  yield 
pure  nitrogen  chloride  by  a second  treatment  with  chlorine  [see  Gatter- 
mann,  Ber.  21  (1888),  751].  A pentachloride  of  nitrogen,  NCI5,  is  not 
known. 

Nitrogen  chloride  is  an  oily,  yellow  liquid,  with  a disagreeable  odor; 
its  specific  gravity  equals  1.65.  It  does  not  solidify  even  at — 40°.  It 
is  an  extremely  dangerous  compound,  as  it  decomposes  by  the  slightest 
contact  with  many  substances,  when  the  temperature  exceeds  90°  and  by 
the  action  of  direct  sunlight.  Its  decomposition  is  accompanied  by  an 
extremely  violent  report.  Its  solution  in  benzene,  ether  and  carbon 
bisulphide  can  be  handled  with  comparatively  little  danger;  but  it 
rapidly  decom])oses  in  them.  Aqueous  ammonia  gradually  decomposes 
it  into  ammonium  chloride  and  nitrogen  : 

NC]3  + 4NH3  = 3NH,C1  + N,. 

It  is  converted  into  ammonium  chloride  and  free  chlorine  by  concen- 
trated hydrochloric  acid : 

NCI3  -f  4TIC1  = NIbCl  -f  3Cb. 

This  reaction  is  directly  opposed  to  that  by  which  nitrogen  chloride  is 
formed. 


34 


INORGANIC  CIIF.MISTRY. 


Nitrogen  cliloride  is  (lccoini)osed  l)y  water  jjartly  into  liypochI(j- 
roiis  acid  [see  Jkr.  27  (1894),  i,  1017,  and  30  (1897),  ii,  1434,  1792; 
III,  2642]. 

'File  formation  and  explosihility  of  nitrogen  cliloride  may  be  illustrated  in  a harmless 
way  as  follows:  Decompose  a saturated  ammonium  chloride  solution  with  the  electric 
current.  Nitrogen  chloride  rising  in  small  drops  from  the  li(iuid  will  sejiaratc  at  the 
])ositive  pole.  Upon  covering  the  surface  of  the  solution  with  a thin  layer  of  turpentine 
oil,  each  drop  will  explode  as  it  comes  in  contact  with  the  latter. 

Nitrogen  Iodide.  Uiion  adding  aininonimn  liydroxide,  ora  mix- 
ture of  ammonium  chloride  and  caustic  soda,  to  a soltition  of  iodine  in 
atjneotis  potassium  iodide,  a brownish-ldack  powder  separates.  Its  com- 
position  closely  approximates  the  formula,  NIII^.  Its  formation  by 
means  of  ammonium  chloride  and  caustic  soda  is  rejiresented  in  the 
equation  : 

4I  + NIbCl  4-  sNaOII  = NIII2  + 3II2O  4-  NaCl  j 2NaI. 

[Com]iare  Raschig,  Ann.  Chem.  230  (1885),  212.] 

When  the  conditions  are  slightly  changed  a very  similar  compound 
separates.  Its  formula  is  N2l3H3(=  Nil,  -f  NI3).  Protracted  washing 
with  water  decomposes  it  into  ammonia  and  nitrogen  tri-iodide,  NI,. 

Nitrogen  Di-iodide  and  Nitrogen  Tri-iodide,  NHIj  andNl,,  are, 
when  dry,  very  explosive.  The  explosihility  may  be  shown  without 
danger  in  the  following  manner:  The  precipitate  is  collected  on  a filter, 
washed  with  water,  the  filter  opened  outand  torn  into  small  pieces,  which 
are  then  allowed  to  dry;  upon  the  slightest  disturbance  these  pieces 
explode  with  a sharp  report. 

Szuhay  claims  that  nitrogen  iodide,  NHI2,  behaves  like  a feeble  acid 
and  that  its  silver  salt,  NAgl2,  explodes  [Ber.  26  (1893),  ii,  1933]. 

Nitrogen  iodide  dissolves  in  dilute  hydrochloric  acid  and  decomposes 
into  ammonia  and  iodine  monochloride  : 

NII2I  + HCl  = NH3  4-  ICl. 

H\  drogen  sulphide  and  sulphurous  acid  convert  it  into  ammonia  and 
hydrogen  iodide. 

'Fhe  nitrogen  iodide  formed  by  digesting  powdered  iodine  with  ammo- 
nia water  manifests  properties  which  are  slightly  different  from  those  of  the 
ordinary  iodide.  It  is  only  stable  in  the  presence  of  ammonia.  It  some- 
times explodes  even  when  moist — if  it  be  washed  with  water,  or  when 
acted  ui)on  by  hydrochloric  acid. 

TJie7'mo-cheniical  Deporfmenf. — Nitrogen  chloride  and  iodide  are  both  strongly  endo- 
thermic comjKHinds  ; considerable  heat  is  absorbed  in  their  production  from  the  elements, 
d'his  ngain  ])resup])oses  conditions  under  which  the  nitrogen  halides  would  be  no  longer 
stable.  ( ’onsc(|uently  their  heat  of  formations  cannot  be  obtained  by  direct  measurement 
but  must  be  calculated  from  other  thermo-chemical  relations.  That  for  nitrogen  chloride  is 


(N,Cl3)  = -41.9  Cal. 


PHOSPHORUS. 


135 


2.  PHOSPHORUS. 

Atom:  P = 3i.o.  Molecule:  P4=  124.0. 

This  element  does  not  occur  free  in  nature,  because  of  its  very  great 
affinity  for  oxygen.  The  phosphates,  especially  calcium  phosphate,  are 
widely  distributed.  By  the  disintegration  of  the  minerals  containing 
phosphates  the  latter  pass  into  the  soil,  are  absorbed  by  plants,  and  these 
are  consumed  by  animals.  Hence  phosphates  occur  in  both  plant 
and  animal  ashes.  The  bones  of  animals  consist  chiefly  of  calcium 
phosphate. 

Brand,  in  Hamburg,  and  Kunkel,  in  Wittenberg  (1670),  obtained 
phosphorus  almost  simultaneously  by  the  ignition  of  evaporated  urine. 
In  1755,  Scheele,  in  Sweden,  showed  that  it  could  be  obtained  from 
bones,  in  which  Gahn  had  found  calcium  phosphate.  Its  name  is  derived 
from  its  power  of  giving  light  in  the  dark — (puxTipopoq,  light-bearer. 

Phosphoru.s  is  prepared  in  France  and  in  England  almost  exclusively  in  accordance 
■with  a method  first  suggested  in  1778  by  Nicolas  and  Pelletier,  which  method  was 
subsequently  improved  by  Fourcroy  and  Vauquelin. 

The  bones  are  burned,  thereby  destroying  all  organic  admixtures,  or  they  are  first 
treated  with  .superheated  steam  to  remove  the  bone  glue.  The  bone  ashes  consist  princi- 
pally of  tertiary  calcium  phosphate  Ca3(P04)2  (see  Phosphoric  Acid).  They  are  dige.sted 
with  two-thirds  of  their  weight  of  concentrated  sulphuric  acid,  when  the  tertiary  phos- 
phate becomes  primary  calcium  phosphate,  and  gypsum  (calcium  sulphate)  is  produced  : 

Ca3(P04)2  + 2H2SO4  = CaH,(POd2  + 2CaS04. 

Tertiary  Primary  Calcium 

calcium  phosphate.  calcium  phosphate.  sulphate. 

The  gypsum,  which  dissolves  with  difficulty  in  water,  is  separated  from  the  readily 
.soluble  primary  phosphate  by  filtration  ; the  solution  is  mixed  with  charcoal,  evaporated 
in  leaden  pans,  and  the  residue  raised  to  a red  heat  in  clay  retorts.  This  expels  water 
from  the  primary  phosphate  and  the  latter  changes  to  calcium  metaphosphate  : 

CaH,(POj2  = Ca(P03b  + 211^0. 

Calcium 

metaphosphate. 

The  ignited  residue  is  then  raised  to  a white  heat.  The  carbon  partly  reduces  the 
metaphosphate  to  ])hosphorus,  by  forming  carbon  monoxide  with  oxygen.  The  resulting 
calcium  oxide  in  like  manner  converts  a portion  of  the  metapho.sphate  into  tertiary  phos- 
phate, which  is  not  further  attacked  by  the  carbon.  The  following  equations  show  the 
changes : 

I.  2Ca(P03).2  + loC  = P,  f loCO  + 2CaO 
II.  Ca(P03).2  4-  2CaO  Ca3(POb2 

3Ca(P03).2  + loC  = P4  -f-  Ca3(P04)2  4 loCO. 

Recently  the  great  heat  of  the  electric  furnace  has  been  apjdied  in  carrying  out  a 
method  suggested  by  Wohler,  but  which  has  not  until  the  pre.sent  been  made  operative  : 
the  tertiary  pho.sphate,  mixed  with  .sand,  is  reduced  by  carbon.  The  silicic  acid  (sand) 
acts  like  the  .sulphuric  acid,  in  that  it  liberates  phosphoric  acid  when  heated  : 

2Ca3(PO^)2  -|-  loC  -(-  6Si02  = 6CaSiO.,  -f-  loCO  4 4B 

Mineral  phosphates  have  been  worked  for  phosphorus  in  a similar  way. 

The  liberated  phosphorus  e.scapes  in  vapor  form,  and  is  collected  and  conden.sed  under 
water  in  receivers  of  peculiar  con.struction.  To  remove  mechanically  admixed  impurities 


T36 


INORGANIC  CHF-MISTRY. 


the  phosphorus  is  aji^ain  distilled  from  retorts,  melted  under  water,  pressed  through  leather 
or  chamois,  or  through  bone  ash,  and  finally  moulded  into  sticks. 

Phosphorus  may  be  obtained  by  the  method  of  Rossel  on  a small  scale,  if  2.5  parts  of 
aluminium  powder,  6 parts  of  sodium  metaphosphate,  and  2 parts  of  silica  are  carefully 
heated  in  a glass  tube  in  a current  of  hydrogen  : 

bNaPOj  -f  loAl  -f  3Si02  = sNa^SiOj  -f  SAl^Oj  -f  6P. 

(Compare  Chemiker-Zeitung,  1898,  237.) 

The  almost  colorless  })hos})horus  obtained  by  distillation  is  a waxy, 
transparent  substance,  of  specific  gravity  1.83  at  10°  C.  At  ordi- 
nary temperatures  it  is  soft  and  tough;  at  o'^  it  becomes  brittle.  It  fuses 
under  water  at  44.4°  and  boils  at  287°.  By  the  action  of  sunlight  it 
becomes  yellow  (hence  called  yellow  phosphorus),  and  is  coated  with  a 
non-transparent,  reddish-white  layer.  Phosphorus  is  insoluble  in  water, 
slightly  soluble  in  alcohol  and  ether,  and  very  readily  soluble  in  carbon 
bisulphide.  It  crystallizes  from  the  latter  solution  in  forms  of  the  isometric 
system  (rhombic  dodecahedra).  When  exposed  to  moist  air,  it  oxidizes  to 
phosphorous  acid  (HgPOg)  and  phosjihoric  acid  (H^POJ  ; the  white  vajiors 
which  arise  contain  ammonium  nitrite  (NH^NOJ,  ozone,  and  hydrogen 
peroxide.  The  odor  of  phosphorus  resembles  that  of  ozone.  It  phosphor- 
esces in  the  dark  on  exposure  to  air.  It  does  this  also  in  other  gases,  but  only 
in  such  as  contain  oxygen.  It  appears  the  phos])horescence  is  influenced  by 
the  formation  and  combustion  of  the  self-inflammable  phosphine,  as  all 
substances  which  destroy  the  latter  prevent  and  put  an  end  to  the  former. 
It  is  noteworthy  that  in  pure  oxygen  the  oxidation  of  phosphorus  begins 
at  27°.  If  the  oxygen  be  diluted  by  removal  over  an  air  pump  or  by  the 
addition  of  neutral  gases,  so  that  its  quantity  is  not  more  than  40  per  cent., 
the  absorption  will  be  very  energetic  at  20°,  but  cease  entirely  at  7°. 

The.  7'ed  phosphorus — discovered  by  v.  Schrotter  in  1845 — possesses 
properties  entirely  different  from  those  of  the  ordinary  variety.  It  is  a 
reddish-brown  powder,  of  specific  gravity  2.19  ; is  not  amorphous  as  for- 
merly believed  but,  according  to  the  investigations  of  Retgers,  it  is 
hexagonal  (Z.  f.  anorg.  Chem.  v (1894),  211).  It  is  insoluble  in  carbon 
bisulphide,  non-phosphorescent,  does  not  alter  in  the  air,  and  is,  indeed, 
very  stable.  While  ordinary  phosphorus  is  very  poisonous,  the  red  variety 
is  perfectly  harmless.  It  does  not  fuse  at  a red  heat,  even  when  subjected 
to  strong  pressure,  and  vaporizes  very  slowly  (even  at  100°),  without 
changing  to  the  yellow  variety;  but  when  it  is  heated  rapidly  (above 
260°)  the  vapors  change  to  those  of  yellow  phosphorus. 

To  prepare  the  red  variety,  yellow  phosphorus  is  heated  for  some  min- 
utes to  300°,  in  closed,  air-tight  iron  vessels;  there  is  a partial  conver- 
sion at  250°.  The  resulting  mass  is  then  treated  with  carbon  bisulphide 
or  sodium  hydroxide,  to  withdraw  the  unaltered,  ordinary  phosphorus. 
If  iodine  be  added  to  the  ordinary  phosphorus,  the  change  will  occur 
below  200°. 

A third  modification — f}iefallic  phosphorus — is  formed  if  the  amorphous  variety  be 
heatcfl  in  a glass  tube,  free  from  air,  to  530°.  Microscopic  needles  then  sublime  into  the 
upper,  less  healed,  portion  of  the  tube.  It  is  more  easily  obtained  if  phos|)horus  is 
heated  with  lead,  in  a closed  tube,  to  a red  lieat.  The  molten  metal  dissolves  the  phos- 


PHOSPHORUS. 


137 


phorus  and,  on  cooling,  the  latter  separates  in  black,  metallic,  shining  crystals  (llittorf ). 
Metallic  phosphorus  possesses  the  specific  gravity  2.34,  vaporizes  with  difficulty,  and  is 
less  active  than  the  red  variety.  Perhaps  it  consists  of  a better  crystallized  form  of  the 
red  variety. 

Two  green  lines  characterize  the  spectrum  of  phosphorus.  On  con- 
ducting hydrogen  over  a small  piece  of  phosphorus,  heated  in  a glass  tube, 
the  escaping  gas  will  burn  with  a bright  green  flame.  When  ordinary 
phosphorus  is  distilled  with  water,  some  passes  over  with  the  steam  and, 
in  the  dark,  phosphoresces.  This  procedure  serves  for  the  detection  of 
phosphorus  in  poisoning  by  this  substance  (Mitscherlich). 

The  density  of  phosphorus  is  referred  to  O2  = 32  and  therefore  its 
molecular  weight  is  124.  As  the  atomic  weight  of  phosphorus  is  31.0, 
it  follows  that  the  molecule  in  the  form  of  vapor  consists  of  4 atoms  : 

124.0.  We  saw  that  the  sulphur  molecule  at  500°  consists  of  6 atoms 
(Sg),  and  at  900°  of  2 atoms  (S2).  Such  a dissociation  does  not,  however, 
occur  with  phosphorus ; even  at  1040°  its  vapor  density  remains  unaltered, 
although  a partial  dissociation  does  take  place  at  a very  intense  heat. 

When  phosphorus  is  burned  in  oxygen  or  in  air,  it  forms  the  pentoxide 
(P2O5).  The  ordinary  variety  inflames  at  40°,  and  also  by  gentle  fric- 
tion ; the  red  is  not  ignited  below  260°.  The  first  will  burn  with  a 
bright  flame  even  under  water.  To  this  end  heat  pieces  of  phosphorus 
in  a flask  with  water,  until  they  melt,  and  conduct  a current  of  oxygen 
through  the  water.  Phosphorus  combines  very  energetically  with  chlorine, . 
bromine,  and  iodine  at  ordinary  temperatures;  on  throwing  a small  piece 
into  a vessel  containing  dry  chlorine  gas  it  at  once  inflames.  The  red 
only  reacts  with  the  halogens  after  applying  heat.  With  most  of  the 
metals  phosphorus  unites  on  warming,  and  precipitates  some  of  them  from 
solutions  of  their  salts.  From  a silver  nitrate  solution,  it  precipitates 
silver  and  silver  phosphide  (AggP)  ; this  solution,  therefore,  answers  as  a 
counterirritant  in  phosphorus  burns. 

The  difference  in  deportment  of  the  yellow  and  the  red  phosphorus  is  fully  accounted 
for  by  the  circumstance,  that  when  the  red  is  produced  from  the  yellow  there  follows  a 
considerable  heat-disengagement : 

P yellow  = P red  +19.2  Cal. 

Hence,  the  red  variety  contains  much  less  energy  than  the  yellow.  In  its  union  with 
other  substances  there  will  always  be  liberated  19.2  Cal.  less,  and  the  reaction  conse- 
f|uently  will  proceed  more  sluggishly  and  with  less  energy.  It  may  be  assumed,  that  in 
the  yellow  phosphorus,  owing  to  the  very  energetic  motions  of  its  atoms  their  mutual 
affinity  acts  in  a much  reduced  form  ; the  atoms  are  loosely  combined  with  one  another 
and  are  therefore  very  reactive.  In  the  pas.sage  of  the  yellow  into  the  red  or  black 
variety  the  energy  of  the  atom  motions  is  converted  into  thermal  energy,  and  the  atoms 
can  attach  themselves  more  firmly  to  one  another  than  before,  the  molecule  is  more  stable, 
the  reactivity  is  diminished.  Formerly  when  heat  was  still  regarded  as  a substance,  such 
relations  as  those  above  were  well  expres.sed  by  the  terms  calorons — bodies  containing 
but  little  heat,  and  calorids — bodies  rich  in  heat. 


12 


INORGANIC  CHRMIS'IRY. 


COMPOUNDS  OF  PHOSPHORUS  WITH  HYDROGEN. 

Plla,  RJI^, 

The  compounds  of  phosphorus  with  hydrogen  can  be  jirejjared  l)y  the 
acticni  of  nascent  hydrogen  upon  ])liosp]u)rus,  as,  for  exainjile,  on  gently 
heating  dilute  suljihuric  acid  with  zinc  and  ])hosphorus  (coinj)arc  ]>.  144). 
I he  usual  course  is  to  heat  yellow  jihosphorus  with  concentrated  potassium 
or  sodium  hydroxide,  when  si)ontaneously  inflammable  i)hosi)hine  escai)es 
and  a salt  of  hypo})hosj)horous  acid  remains  in  solution. 

The  lil)erate(l  jras  mixed  with  air  in  a closed  ve.ssel  explodes  violently  ; hence,  to  make 
it,  proceed  as  follows  : hill  a small  glass  flask  almost  full  of  a(jueous  pota.ssium  hydroxide, 


Fig.  55. 


add  a few  pieces  of  phosphorus,  and  heat  over  a lamp  (Fig.  55).  When  the  liberation  of 
gas  commences,  and  the  air  in  the  neck  of  the  flask  has  been  expelled,  close  the  .same  with 
the  cork  of  the  delivery  tube,  the  other  end  of  which  dips  under  warm  water,  to  prevent 
any  obstruction  arising  in  it  from  phosphorus  that  may  be  carried  over  and  solidify  by 
cooling.  Kach  bubble  rising  from  the  liquid  inflames  in  the  air,  and  forms  white  cloud- 
rings  which  ascend. 

I he  gas  tints  ])roduced  consists  of  gaseous  plios|)hine  (PH.j)  and  hydro- 
gen, with  whicli  is  mixed  a small  (piantity  of  a licpiid  substance 
whose  presence  imparts  the  s|)ontaneous  inflammability  to  the  gas.  On 
ctmducting  the  latter  through  a cooled  tube  the  P^H^  is  condensed  to  a 
lifluid,  and  the  escaping  gas  no  longer  inflames  sitontaneously.  The 


COMPOUNDS  OF  PHOSPHORUS  WITH  HYDROOFN. 


139 


licpiid  compound  may  be  isolated  in  a similar  manner  if  tlie  gas  is  con- 
ducted through  alcohol  or  ether,  which  will  absorb  the  cum})ound 

Liquid  Phosphine,  P2H4>  separated  from  the  gas  by  cooling,  is  a colorless,  strongly 
refracting  liquid  of  specific  gravity  1.012,  insoluble  in  water,  and  boiling  at  57°  [Gat- 
termann  and  llausknecht,  Ber.  23  (1890),  1174].  It  inflames  spontaneously  in  the  air, 
and  burns  with  great  brilliancy  to  phosphorus  pentoxide  and  water.  Its  jjresence  in  com- 
bustible gases,  such  as  hydrogen,  marsh  gas,  and  gaseous  phosj)hine,  gives  them  their 
spontaneous  inflammability.  In  contact  with  .some  compounds,  like  carbon  and  sulphur, 
and  by  the  action  of  sunlight  and  of  strong  hydrochloric  acid,  it  decomposes  into  gaseous 
and  solid  phosphine  f 

5P,H,  = 6PH3-f  lyi,. 

Solid  Phosphine,  P4II2  (?)  is  a yellow  powder,  which  inflames  at  160°  or  by  a blow. 
It  is  produced  in  the  decomposition  of  calcium  phosphide  by  hydrochloric  acid. 

Gaseous  Phosphine,  PH3,  may  be  formed,  together  with  the  liquid 
and  solid  variety,  in  addition  to  the  manner  previously  described,  by 
the  action  of  water  or  hydrochloric  acid  upon  the  calcium  phosphides, 
Ca3P2  and  Ca2P2 : 

Ca3P2  -f  6HC1  = 3CaCl2  -j-  2PH3 ; 

Ca2P2  + 4IICI  = 2CaCl2  + P2H4 ; 

5P2ll4=6PH3  -fP4ll2. 

Further,  by  the  ignition  of  phosphorous  and  hypophosphorous  acids : 

4H3PO3  = PH3  + 3H3PO, 

Phosphorous  Phosphoric 

acid.  acid. 

and  by  decomposing  phosphonium  iodide  with  caustic  potash  (p.  140), 

It  is  a colorless  gas,  with  a disagreeable,  garlic-like  odor,  and  is 
slightly  soluble  in  alcohol.  Its  density  is  34  03  referred  to  O2  = 32, 
or  1.176  (air=  i).  It  solidifies  at  — 133- 5°  and  boils  at  about  — 85°. 
When  i)ure  (free  from  P2HJ  it  ignites  at  100°.  Oxidizing  agents  con- 
vert it  again  into  the  spontaneously  inflammable  variety,  owing  to  the 
production  of  PgH^.  It  is  extremely  poisonous.  Phosphine  is  decom- 
posed into  phosiihorus  and  hydrogen  when  it  is  heated,  or  if  it  is  exposed 
to  the  action  of  the  electric  spark.  When  ignited  in  the  air  it  burns 
with  a brightly  luminous  flame,  disseminating  at  the  same  time  a white 
cloud  of  phosphorus  pentoxide  (P2O5) : 

2PII3  + 4O2  = 3H2O  -f  P2O5. 

When  mixed  with  chlorine  it  explodes  violently,  with  production  of 
phosphorus  trichloride  and  hydrogen  chloride: 

PH3  4-  3CI2  - PCI3  + 3HCI. 

Like  ammonia,  phosphine  jiossesses  faint  alkaline  properties,  and  com- 
bines with  hydrogen  iodide  and  bromide  to  yield  compounds  similar  to 
ammonium  chloride  : 

PII3  -f  HI  ==  PHJ. 


It  combines  with  hydrogen  chloride  at  from  — 30*^  to — 35°,  or,  at  ordi- 


140 


INORGANIC  CHEMISTRY. 


nary  tem])eratiires,  under  a jiressure  of  20  atmosplicrcs.  'Die  groiiji  PII^, 
figuring  in  the  role  of  a metal  in  tliese  compounds,  is  analogous  to 
ammonium  (p.  129),  and  is  termed  pJiosphoniuin. 

Phosphonium  Iodide,  PHJ.  Jt  is  best  prepared  by  the  decomposi- 
tion of  phosphorus  di-iodide  (ly^,  p.  142),  by  a slight  (quantity  of  water,  or 
by  adding  yellow  jihosphonis  (10  jiarts),  and,  after  some  hours,  iodine  (2 
jiarts),  to  a saturated  solution  of  hydriodic  acid  (22  jiarts).  The  bhpiid 
becomes  a solid  mass,  consisting  of  ])hosphonium  iodide  and  phosjihoroiis 
acid.  Phosphonium  iodide  sublimes  in  colorless,  shining,  cube-like 
rhombohedra  and  by  this  means  may  be  obtained  jiiire.  It  fumes  in  the 
air,  and,  with  water,  decomposes  into  phosphine  and  hydrogen  iodide. 
When  decomposed  by  potassium  hydroxide  it  yields  pure  hydrogen 
phosjihide,  which  is  not  spontaneously  inflammable  : 

PI  1,1  + icon  =3  KI  + PH3  4-  II/). 

Phosphine  is  a feebly  exothermic  compound  : 

P yellow  + 311  = PII3  + 1 1.6  Cal. 


MOLECULAR  FORMULA  OF  PHOSPHINE.  ATOMIC  WEIGHT  OF 
PHOSPHORUS. 

The  analysis  of  phosphine  shows  that  it  consists  of  i.oi  parts  of  hydrogen  and  10.33 
parts  of  phosphorus.  Were  its  molecular  formula  PH,  the  atomic  weight  of  phosphorus 
would  be  10.33.  The  great  analogy  of  phosphine  to  ammonia,  and  that  of  all  the 
phosphorus  compounds  to  those  of  nitrogen,  argues,  however,  for  the  formula  PII3. 
The  atomic  weight  of  phosphorus,  therefore,  is  31.00  (=  3 X ^0.33),  and  the  molecu- 
lar weight  of  the  phosphine  is  34.03  : 

113=  3.03 
P =31.00 


PH3  = 34.03 

This  view  is  confirmed  by  the  density.  Direct  experiment  confirms  this.  Further, 
from  the  formula  PH3  it  follows  that  3 volumes  of  hydrogen  are  present  in  2 volumes  of 
the  gas : 

2PH3  contain  3H2, 

2 vols.  3 vols. 

or  in  I volume  there  are  volumes  of  hydrogen.  On  decomposing  the  gas  in  a 
eudiometer,  by  means  of  electric  sparks,  it  will  be  found  that  the  volume  increases  one-J 
half ; the  gas  consists,  then,  of  pure  hydrogen,  while  phosphorus  .separates  in  a solid 
condition.  As  the  phosphorus  molecule  in  the  gaseous  condition  is  composed  of  four 
atoms  (p.  137),  the  jdiosphorus  (62.00  parts)  sej)arated  from  2 volumes  of  phosphine 
will  fill  X volume  when  in  the  form  of  vapor;  hence  in  2 volumes  of  phosphine  there 
are  present  3 volumes  of  hydrogen  and  X volume  of  phosphorus  vapor. 

Or,  written  molecularly  : 

4Pll3  = P4  -f  611.2. 

1 vol,  I vol.  6 vols. 


COMPOUNDS  OF  PHOSPHORUS  WITH  THE  HALOGENS. 


I4I 


COMPOUNDS  OF  PHOSPHORUS  WITH  THE 
HALOGENS. 

Phosphorus  combines  directly  with  the  halogens  to  yield  compounds 
of  the  types  PX3  and  PX^,  in  which  X indicates  an  halogen  atom. 

Phosphorus  Trichloride — Phosphorous  Chloride — PCI3.  Conduct 
dry  chlorine  gas  over  phosphorus  gently  heated  in  the  retort  C (Fig.  49, 
p.  in).  The  phosphorus  ignites  in  the  stream  of  gas,  and  distils  over  as 
trichloride,  which  is  collected  in  the  receiver  Z>,  and  condensed.  The 
product  is  purified  by  a second  distillation.  It  is  a colorless  liquid,  con- 
gealing at  — 112°,  boiling  at  76°,  and  has  a sharp,  penetrating  odor. 
Its  specific  gravity  equals  1.613  at  0°.  It  fumes  strongly  in  the  air,  and 
is  decomposed  by  moisture  into  phosphorous  and  hydrochloric  acids  : 

PCI3  + 3H2O  = H3PO3  + 3HCI. 

The  vapor  density  of  the  trichloride  is  137.35,  corresponding  to  the 
molecular  formula  PCI3. 

Phosphorus  Pentachloride — Phosphoric  Chloride — PCI5.  This  is 
produced  by  the  action  of  an  excess  of  chlorine  upon  the  liquid  trichlo- 
ride. It  is  a solid,  crystalline,  yellowish-white  compound.  It  fumes 
strongly  in  the  air  and  sublimes  without  melting  when  heated.  It  at  the 
same  time  sustains  a partial  decomposition  into  trichloride  and  chlorine. 
Under  increased  pressure  it  melts  at  148°. 

In  an  atmosphere  of  phosphorus  trichloride  the  vapor  density  of  the  pentachloride 
has  been  found  to  be  208.25,  corresponding  to  the  molecular  formula  PCI5  = 208.25. 
At  increased  temperatures  the  vapor  density  steadily  diminishes,  and  a gradual  de- 
composition occurs — dissociation  of  the  molecules  of  the  pentachloride  (PCI5)  into 
the  molecules  of  the  trichloride  (PCI3),  and  chlorine  (Cb).  The  decomposition  temper- 
ature, /.<?.,  the  temperature  at  which  the  decomposition  is  half  finished,  lies  at  about 
200°  C.  The  dissociation  is  complete  at  300°,  and  equals  104,  the  vapor  density;  i.  e., 
the  vapor  then  fills  a volume  twice  as  large  as  at  a lower  temperature.  The  breaking 
down  of  the  pentachloride  into  the  trichloride  and  chlorine,  whereby  2n  molecules  result 
from  n molecules,  explains  this  : 

PCb  = PCI3  4-  Cb. 

I vol.  I vol.  I vol. 

That  such  a decomposition  of  the  penta-  into  trichloride  and  chlorine  does  really  occur, 
is  y^roved,  among  other  things,  by  the  originally  colorless  vapor  gradually  assuming  the 
yellow  color  of  chlorine  as  the  temperature  rises.  The  decomposition  j:)roducts — the 
trichloride  and  chlorine — may  be  separated  from  each  other  by  diffusion. 

Phosphorus  pentachloride  acts  very  energetically  with  water ; it  yields, 
depending  upon  the  quantities  of  the  reacting  substances,  phosphorus 
oxychloride,  POCI3,  metaphosphoric  acid,  HPO3,  or  orthophosphoric 
acid,  H3pO^,  and  hydrochloric  acid  : 

PC'b  + 11,0  = POCI3  -f  2IICI 
PCb  + 3H,0  = HPO,  + 5IICI 
PCb  4-  41130  = I^PO^  4-  5HCI. 

(See  Phosphoric  Acid.) 


142 


INORGANIC  CIIKMISTUY. 


"I'he  ])ioiiiiiie  and  iodine  })h()sphonis  compounds  are  jierfeclly  analogous 
to  the  chlorine  derivatives.  They  are  obtained  by  uniting  ihe  constitu- 
ents in  the  ])roi)ortions  by  weight  expressed  by  their  formulas.  As  the 
union  is  exceedingly  energetic,  it  is  best  to  ])roceed  as  follows:  Dissolve 
the  phosphorus  in  carbon  bisulj)hide,  gradually  add  the  calculated  amount 
of  bromine  or  iodine,  and  then  distil  off  the  volatile  solvent. 


Phosphorus  Tribromide,  PHr.,,  is  a colorless  liquid,  boiling  at  175°,  and  having  a 
specific  gravity  of  2.9  at  0°.  'I'lie  pentabromide,  PP>r^,  formeil  by  the  gradual  acUlition 
of  2Br  to  PBr-j,  is  a yellow,  crystalline  substance,  which  melts  when  heated,  and  breaks 
down  into  the  tribromide  and  bromine.  Water  decomposes  both  compounds,  as  it  does 
the  corresponding  chlorides. 

Phosphorus  Chlorbromide,  PCbIjrj,  is  ])roduced  by  the  union  of  PCI,  with  Br^  in 
the  cold.  It  is  a yellowish-red  mass,  which  decomjKises  at  35°  into  PCI.,  and  Br.^. 

Phosphorus  Tri-iodide,  PI,,  forms  red  crystals,  melting  at  55°  and  distils,  with  par- 
tial decom|)osition,  at  a higher  temperature.  'I'he  so-called  phosphorus  iodide,  P.^b 
(corresiK)nding  to  P.^H,),  crystallizes  in  beautiful  orange-red  needles  or  prisms,  ami 
fuses  at  110°.  Its  vapor  density  at  265°  and  90.7  mm.  })ressure  ecpials  569,  correspond- 
ing to  the  molecular  weight  P^b-  A little  water  decomposes  it  into  phosjihorous  acid, 
phosi)hine,  and  hydriodic  acid.  The  last  two  bodies  then  form  phosphoniiini  iodide, 

PIbl  (p.  140). 

The  recently  di.scovered  Phosphorus  Pentafluoride,  PFl^,  is  interesting  (Thorpe). 
It  results  upon  heating  phosphorus  trichloride  or  pentachloride  with  arsenic  trifluoride, 
AsFl,  : 

3PCb  + 5ASFI3  = 3PFb  4 SAsCl,. 

It  is  a colorless  gas  which  fumes  in  moist  air  and  is  decomposed  by  water  into  phos- 
phoric acid  and  hydrogen  fluoride.  Its  density  corresponds  to  the  molecular  formula 
PFl^  =:  125.9.  It  may  be  liquefied  at  16°  under  a pressure  of  46  atmospheres,  and 
solidifies  when  the  pressure  is  removed. 

It  is  rather  remarkable  that  although  phosphorus  pentiodide  could  not  be  obtained, 
the  stability  of  the  compounds,  PBr,,  PCl^,  PFl-,  gradually  increases  with  the  diminution 
of  the  atomic  weight  of  the  combined  halogens.  Phosphorus  pentafluoride  can  be  gasi- 
fied without  decomposition. 


Thermo-chemical  Deportfueiit. — While  the  halogen  derivatives  of  nitrogen  (like  tho.se 
of  oxygen)  are  strongly  endothermic,  are  produced  wdth  the  absorption  of  much  heat,  and 
are,  in  consequence,  readily  exploded  (p.  134),  those  of  phosphorus  are  exothermic.  The 
heat  disengaged  in  the  union  of  yellow  phosphorus  and  chlorine  (p.  137)  corresponds  to 
the  following  symbols  : 

(P,Cl3)  = 75.5;  (P,Cb)  = io5. 

In  this  we  observe  a transition  to  the  halogen  derivatives  of  the  metals,  all  of  whicli 
are  exothermic.  In  accordance  with  this  we  find  that  the  heat  of  formation  of  the  bro- 
mides and  iodides  diniini.shes  in  regular  succession  : 

(DCb)  = 75-5 ; (DB‘3)  = 45 ; (PJ3)  = io-9- 

'riie  great  reactivity  of  all  the.se  derivatives  with  water  is  fully  ex]>lained  by  the  large 
amount  of  heat  .set  free  at  the  .same  time — the  decomposition  of  phosphorus  tri- 
chloride with  .so  much  water  that  upon  diluting  the  resulting  .solution,  no  further  heat 
evolution  takes  place  : 

PCd,  -f  A(i  — IbPOsAq  4-  3lIClAq  . . . 4-65  Cal. 


ARSENIC. 


143 


3.  ARSENIC. 

Atom  : As  = 75.  Molecule  : As^  = 300, 

Arsenic  is  a perfect  analogue  of  phosphorus,  but  possesses  a somewhat 
metallic  character.  In  its  free  state  it  is  similar  to  metals. 

Arsenic  is  found  free  in  nature,  although  it  occurs  more  frequently  in 
combination  with  sulphur  (realgar,  As.^S.^,  orpiment,  As^Sg),  with  oxygen 
(arsenolite,  As.^03),  and  with  metals  (mispickel,  FeAsS,  cobaltite,  CoAsS). 
To  prepare  it,  heat  mispickel  with  iron,  and  free  arsenic  will  sublime, 
while  iron  sulphide  remains.  Or,  in  the  customary  way  of  isolating  metals 
from  their  oxides,  heat  the  trioxide  (arsenolite)  with  charcoal : 

2AS2O3  6C  = As^  -[-  6CO. 

Arsenic  appears  in  two  modifications.  Crystallized  (hexagonal)  arsenic 
is  obtained  by  the  sublimation  of  ordinary  arsenic.  It  forms  a gray- 
white,  more  or  less  metallic,  crystalline  mass,  but  may  be  changed  into 
acute  rhombic  octahedra.  Its  specific  gravity  equals  5.73.  It  is  brittle, 
and  may  be  pulverized  without  difficulty. 

The  amorphous  is  microcrystalline  (according  to  Retgers)  and 

probably  isometric.  It  is  formed  along  with  the  first  variety  when  arsenic 
is  sublimed  in  a glass  tube  in  a current  of  hydrogen  (Bettendorff ) and 
also  upon  heating  arsine.  It  is  black,  with  little  luster,  and  possesses  the 
specific  gravity  4.71.  When  heated  to  360°  it  sets  heat  free  and  reverts 
to  the  crystalline  variety. 

Away  from  air  contact,  and  at  the  ordinary  pressure,  arsenic  vapor- 
izes at  a dark-red  heat  (about  450°)  without  previously  melting  ; it  will, 
however,  melt  if  heated  under  great  pressure  in  a sealed  tube.  Its  vapor 
possesses  a lemon-yellow  color.  The  vapor  density  corresponds  to  the 
molecular  weight.  As  its  atomic  weight  equals  75,  it  follows  that  the 
molecule  in  the  gaseous  state  consists,  like  that  of  phosphorus,  of  four 
atoms.  At  a white  heat  (about  1700°)  the  density  falls  to  one-half, 
which  is  due  to  the  fact  that  then  the  arsenic  vapor  consists  largely  of 
diatomic  molecules  (Asj  = 150). 

Arsenic  does  not  change  in  dry  air.  When  heated  in  the  air  it  inflames 
at  180°  and  burns  with  a blue-colored  flame,  disseminating  the  garlic-like 
odor  of  arsenic  trioxide  (As^Og).  It  combines  directly  with  most  ele- 
ments. Powdered  arsenic  will  inflame  when  thrown  into  chlorine  gas. 
It  yields  arsenides  with  the  metals. 

It  is  remarkable  that  arsenic,  belonging  to  the  nitrogen  group  and  generally  forming 
compounds  which  in  constitution  are  quite  different  from  those  of  sulphur,  should  be 
analogous  to  the  latter  in  its  metallic  combinations.  Thus  the  sulphides  and  arsenides 
have  similar  formulas,  are  isomorphous,  and  in  them  sulphur  and  arsenic  can  mutually 
replace  each  other  in  atomic  ratios,  e.  g.  : 

FeS^,  FeAs.^  and  Fe(SAs). 


144 


INORGANIC  CUKMISTRY. 


COMPOUNDS  OF  ARSENIC  WITH  HYDROGEN. 

Arsine,  AsH.,  = 78.03.  Like  nitrogen  and  i)h()sj)liorus,  arsenic  fiir- 
nishes  a gaseous  compound  containing  three  atoms  of  hydrogen.  It  is  ol)- 
tained  i)ure  by  the  action  of  dilute  sulphuric  acid  or  hydrochloric  acid 
upon  an  alloy  of  zinc  and  arsenic: 

As2Zn3  -f  6I1C1  3ZnC1.3  -f-  2ASII3. 

It  also  results  in  the  action  of  nascent  hydrogen  (zinc  and  suliihuric  acid), 
upon  many  arsenic  compounds,  as,  c.  g.^  the  trioxide  : 

AS3O3  ^-6U,  = As,U,-\  3II./). 

Arsine,  discovered  by  Scheele  in  1755,  is  a colorless  gas,  of  strong, 
garlicky  odor,  and  extremely  poisonous  action  ; it  may  be  condensed  to 


a liquid  and  even  to  a solid  by  cold.  It  melts  at  — ii3-5°  and  boils  at 
— 55°.  Its  density  equals  78.03  (O2  = 32)  or  2. 69  (air  = i).  It  burns 
with  a bluish-white  flame  when  ignited,  and  evolves  white  fumes  of 
arsenic  trioxide : 

2ASII3  + 3O2  = AS2O3  -f  3H.p. 

It  is  decomj)osed  at  a dull-red  heat  or  by  the  electric  spark  into 
arsenic  and  hydrogen.  On  conducting  the  gas  through  a heated  tube  the 
arsenic  de|)osits  behind  the  heated  part  as  a metallic  coating  (arse/iic 
mirror).  On  holding  a cold  object,  e.  g.,  a i)iece  of  porcelain,  in  the 
ll.'ime  of  the  gas,  the  arsenic  forms  a black  deposit  (arsenic  sj)ots).  In  its 
rhemical  behavior  arsine  is  very  similar  to  phosi)hine;  its  basic  proper- 
ties are  very  slight,  and  it  does  not  furnish  any  derivatives  with  the 
halogens. 


COMPOUNDS  OF  ARSENIC  WITH  THE  HALOGENS. 


145 


According  to  analysis,  arsine  consists  of  1. 01  parts  by  weight  of  hydrogen  and  25  parts 
of  arsenic.  If,  because  of  its  analogy  to  phosphine,  we  ascribe  it  the  formula  AsH^, 
then  the  atomic  weight  of  arsenic  would  be  75  (=  3 X 25)  and  the  molecular  weight  of 
AsHg  would  equal  78.03.  A determination  of  the  density  confirms  this.  The  formula, 
too,  shows  that  3 volumes  of  hydrogen  are  present  in  2 volumes  of  ASH3  : 

2ASH3  contain  3H2. 

2 vols.  3 vols. 

We  can  satisfy  ourselves  of  this  by  decomposing  the  gas  by  electricity  in  a eudiometer 
(see  p.  120). 

Marsh’s  Method  for  the  Detection  of  Arsenic. — The  method  described 
for  the  preparation  of  arsine  and  the  ease  with  which  it  is  recognized 
make  it  possible  to  detect  arsenic  in  its  compounds  with  great  certainty. 
This  is  a very  important  task  because  of  the  ready  accessibility  and  fre- 
quent use  of  poisonous  arsenic  derivatives.  Hydrogen  is  generated  in  a 
flask  (Fig.  56,  a)  by  the  interaction  of  dilute  sulphuric  acid  and  zinc, 
when  the  material  presumed  to  contain  arsenic  is  introduced  through  the 
funnel  tube  b.  The  gas  evolved,  a mixture  of  hydrogen  and  arsine,  is 
then  conducted  through  a calcium  chloride  drying  tube  (c)  and  escapes 
through  the  tube  of  hard  glass  contracted  at  several  points  (d).  Upon 
igniting  the  escaping  hydrogen  (after  all  the  air  has  been  previously  expelled 
from  the  vessel,  as  otherwise  oxyhydrogen  gas  will  be  present)  it  will 
burn  with  a bluish-white  flame,  if  arsenic  be  present,  and  at  the  same  time 
disseminate  a white  vapor.  The  dark  arsenic  spots  are  obtained  by  hold- 
ing a cold  porcelain  dish  in  the  flame.  If  the  tubed  be  heated  (as  shown 
in  Fig.  56),  an  arsenic  mirror  will  be  formed  upon  the  adjacent  contrac- 
tion. The  slightest  traces  of  arsenic  may  be  detected  by  this  method. 

Besides  the  ordinary  arsine,  AsHg,  we  might  expect  the  existence  of  As-^H^  and 
AS4H2,  corresponding  to  the  liquid  and  solid  phosphines  (P2fl4  and  The  first  is 

not  known  ; its  derivatives  exist,  and  contain  hydrocarbon  groups  instead  of  hydrogen. 
An  example  of  this  class  is  cacodyl,  As2(CH3)^  = (CH3)2As-As(CH3)2.  Nitrogen 
affords  similar  compounds — fCH3)2N-NH2  and  (CH3)NH-NH2,  derived  from  diamide 
or  hydrazine,  N2H^  = H2N-NH2  (p.  131). 

The  solid  arsine,  As^H2,  is  obtained  by  the  action  of  nascent  hydrogen  upon  arsenic 
compounds  in  the  presence  of  nitric  acid.  It  forms  a reddish-browm  powder,  which 
decomposes  when  heated.  Retgers  considers  that  the  brown  color  of  the  arsenic  spots 
and  mirror  is  due  to  solid  arsine,  soluble  in  caustic  potash. 


COMPOUNDS  OF  ARSENIC  WITH  THE  HALOGENS. 

These  are  perfectly  analogous  to  the  corresponding  phosphorus  com- 
pounds, and  are  the  result  of  the  direct  union  of  their  constituents.  The 
fluoride,  in  union  with  potassium  fluoride,  is  the  only  known  representa- 
tive of  the  compounds  corresponding  to  the  formula  AsX^  (see  p.  141). 
The  metallic  character  of  arsenic  is  shown  by  the  fact  that  arsenic  chloride, 
like  other  metallic  chlorides,  may  be  obtained  by  the  action  of  hydro- 
chloric acid  upon  the  oxide  : 

AS2O3  -f  6HC1  = 2ASCI3  -f  3H2O. 


3 


146 


INORGANIC  CHEMISTRY. 


Arstnic  chloride  is  evolved  when  a solution  of  the  trioxide  is  boiled 
with  concentrated  hydrochloric  acid. 

Arsenic  Trichloride,  AsCl,,  is  a colorless,  oily  licjuid,  fuming  in 
the  air,  and  having  a S|)ecific  gravity  of  2.2.  Jt  solidifies  at  lower  tem- 
peratures, melts  at — 18°  and  boils  at  130°.  Its  vapor  density  corresponds 
to  the  molecular  formula  AsCl,  = 181.35.  dissolves  in  a small  quan- 
tity of  water  without  change,  while  much  water  converts  it  into  the  oxide 
and  hydrochloric  acid 

2ASCI3  + 3II2O  = As/J,  -f  6IIC1. 

Arsenic  Tribromide,  AsBr,,  is  a white  crystalline  mass,  melting  at 
20°,  and  boilingat  220°.  The  Tri-iodide,  Asl^,  forms  red  crystals,  which 
melt  at  146”.  The  Trifluoride,  ASFI3,  is  a licpiid,  fuming  strongly  in  the 
air;  it  boils  at  63°.  It  results  in  the  distillation  of  the  trichloride  or  tri- 
oxide with  calcium  fluoride  and  sulphuric  acid. 


4.  ANTIMONY. 

Atom  : Sb  = 120. 

The  metallic  character  exhibited  by  arsenic  becomes  more  distinct  with 
antimony,  which  at  the  same  time  retains  its  complete  analogy  to  the 
metalloidal  elements,  arsenic  and  phosphorus.  Antimony  is  a perfect 
metal  so  far  as  its  physical  properties  are  concerned. 

It  (stibium)  occurs  in  nature  chiefly  in  union  with  sulphur,  as  stibnite, 
^^>283  (Japan,  Hungary),  and  with  sulphur  and  metals  in  many  ores.  It 
is  almost  always  accompanied  by  arsenic.  To  prepare  antimony,  stibnite 
is  roasted  in  a furnace,  i.  e.,  heated  with  air  access,  whereby  the  sulphur 
burns,  and  antimony  trioxide  remains: 

28^83  -f  9O2  = 28b203  -f-  68O2. 

The  residual  oxide  is  ignited  with  carbon,  which  reduces  it  to  metal 
(general  procedure  for  the  separation  of  metals).  Antimony  may  also 
be  obtained  by  heating  its  sulphide  with  iron,  which  combines  with  the 
sulphur : 

SbaSa  + 3re  = 28b  + 3FeS. 

The  resulting  commercial  crude  antimony  is  further  purified  in  the 
laboratory  by  fusing  it  with  niter,  whereby  the  admixed  arsenic,  sulphur, 
and  lead  are  removed.  Chemically  pure  antimony  is  obtained  by  reduc- 
ing the  pure  oxide. 

It  is  a silver-white  and  very  brilliant  metal,  of  leafy  crystalline  struc- 
ture; specific  gravity  6.71.  Like  arsenic  it  crystallizes  in  rhombohedra, 
is  very  brittle,  and  may  be  easily  broken.  It  fuses  at  430^^,  and  vaporizes 
between  1500  to  1700°.  The  density  of  its  vapor  shows  that  the  anti- 
mony molecule,  unlike  that  of  phosphorus,  consists  not  of  four  atoms,  but 
rather,  like  arsenic  at  the  same  temperature,  of  two  atoms  [Riltz,  Z.  f. 
]fliys.  Chem.  ig  (1896),  385].  It  is  not  altered  in  the  air  at  ordinary 
temperatures,  but  when  heated  it  burns  with  a blue  flame,  yielding  white 


COMPOUNDS  OF  ANTIMONY  WITH  THE  HALOGENS. 


147 


vapors  of  antimony  oxide,  SbjO,.  Like  phosphorus  and  arsenic  it  com- 
bines directly  with  the  halogens;  powdered  antimony  inflames  in  chlo- 
rine gas.  It  is  insoluble  in  hydrochloric  acid ; nitric  acid  oxidizes  it 
(depending  upon  its  strength  and  the  temperature)  to  antimony  oxide  or 
antimonic  acid. 

Hydrogen  Antimonide — SHbine — SbHg,  is  produced  like  arsine, 
and  is  very  similar  to  the  latter.  It  has  thus  far  only  been  obtained 
mixed  with  hydrogen.  It  is  a colorless  gas  of  peculiar  odor,  and  when 
ignited  burns  with  a greenish-white  flame,  disseminating  white  vapors  of 
antimony  oxide,  SbgOg.  A red  heat  decomposes  it  into  antimony  and 
hydrogen.  In  Marsh’s  apparatus  (Fig.  56,  p.  144)  it  forms  an  antimony 
mirror  and  spots.  The  mirror  is  distinguished  from  that  of  arsenic  by 
its  black  color,  lack  of  luster,  its  insolubility  in  a solution  of  sodium 
hypochlorite  (NaClO),  and  by  its  slight  volatility  in  a current  of 
hydrogen. 

Arsine  is  also  liberated  from  alkaline  solutions  in  which  hydrogen  is 
being  generated,  e.  g.,  caustic  potash  and  zinc.  This  is  not  the  case 
with  stibine. 


COMPOUNDS  OF  ANTIMONY  WITH  THE 
HALOGENS. 

Antimonous  Chloride — Antimony  Trichloride — SbClg,  results  from 
the  action  of  chlorine  upon  the  metal  or  its  sulphide  (SbgSg) ; better  by 
the  solution  of  the  oxide  or  sulphide  in  strong  hydrochloric  acid  : 

SbgSj  -f  6HC1  = 2SbCl3  -f  3H2S. 

This  solution  is  evaporated  to  dryness  and  the  residue  distilled. 

It  is  a colorless,  crystalline,  soft  mass  {^Butyruni  antimonii),  melting  at 
73°  and  boiling  at  223°.  Its  vapor  density  corresponds  to  the  molecular 
formula,  SbClg  — 226.35.  the  air  it  attracts  water  and  deliquesces. 
It  dissolves  unchanged  in  water  acidified  with  hydrochloric  acid.  Much 
water  decomposes  it;  the  solution  becomes  turbid  and  a white  powder — 
powder  of  Algaroth  (so  named  in  honor  of  an  Italian  physician,  Victor 
Algarotus,  who  used  it  as  a medicine) — separates : 

SbClg  + H2O  = SbOCl  4-  2HCI. 

The  composition  of  this  powder  varies  with  the  conditions  under  which 
it  is  formed,  but  generally  corresponds  to  the  formula  2(SbOCl)  . SbgOg. 
Pure  Antimony  Oxychloride,  SbOCl,  obtained  by  heating  antimony 
trichloride  with  alcohol,  occurs  in  colorless  crystals  and  is  further  decom- 
posed by  water  into  basic  oxychlorides. 

While  the  metallic  chlorides  are  not  decomposed  by  water  at  ordinary 
temperatures,  the  ready  decomposition  of  the  halogen  derivatives  of 
antimony  shows  that  this  element  still  possesses  a partial  metalloidal 
character. 


148 


INORGANIC  CHEMISTRY. 


Antimonic  Chloride — Antimoiiy  Pentachloride — SbCl^,  results  from 
the  action  of  an  excess  of  chlorine  ui)on  antimony  or  the  trichloride.  It 
is  a yellowish  licjuid  which  fumes  in  the  air,  becomes  crystalline  when 
cold  and  melts  at  — 6°.  Heat  partly  decomjioses  it,  like  jihosphorus 
pentachloride,  into  antimonous  chloride  and  chlorine  : 

SbCl^  = SbCf,  4-  Clj. 

I vol.  I vol.  I vol. 

It  may  be  gasified  without  decomposition  at  218°  and  58  mm.  pressure. 
Its  density  corresponds  to  the  formula  SbCl^. 

Water  converts  it  into  pyroantimonic  acid  (H^SbjO^),  and  hydrochloric 
acid.  It  combines  with  one  molecule  of  water,  forming  the  crystallizable 
hydrate  SbCl^ . 11./^,  melting  at  about  90°,  and  with  four  molecules  to  a 
hard,  crystalline  mass,  SbCl5.4H20  [Anschutz  and  Evans,  Ann.  Chem. 
239  (1887),  285].  The  metallic  nature  of  antimony  shows  itself  in  the 
formation  of  these  hydrates.  Neither  of  the  two  chlorides  of  non-metallic 
phosphorus  forms  a hydrate ; both  are  immediately  decomposed  by  a 
little  water  with  the  production  of  hydrochloric  acid. 

Antimony  Tribromide,  SbBrg,  is  a white,  crystalline  substance,  melting  at  93°  and 
distilling  at  275°.  The  Tri-iodide,  Sb^,,  is  a red  compound,  crystallizing  in  three  dis- 
tinct forms.  It  melts  at  166°  and  boils  at  400°.  The  Pentiodide,  Sblj,  is  dark  brown 
in  color  and  melts  at  about  78°. 

Antimony  Trifluoride,  SbFlg,  obtained  by  the  solution  of  antimony  oxide  in  hydro- 
fluoric acid,  crystallizes  in  colorless  rhombic  pyramids.  It  deliquesces  on  exposure  to  the 
air.  It  is  not  decomposed  by  cold  water.  Its  compound  with  ammonium  sulphate  is 
used  as  a mordant  in  dyeing. 

Antimony  Pentafluoride,  SbFlg,  is  a gummy  mass.  It  forms  well-crystallizing 
double  salts  with  the  salts  of  organic  bases. 

We  must  also  include  Bismuth,  Bi  = 208.5,  the  group  of  nitro- 
gen, phosphorus,  arsenic,  and  antimony.  Its  halogen  derivatives  resem- 
ble those  of  arsenic  and  antimony  in  many  respects,  e.  g.,  BiClg,  Bilj, 
BiOCl.  Its  metallic  character,  however,  considerably  exceeds  its  metal- 
loidal.  Thus,  it  does  not  unite  with  hydrogen,  and  bismuth  oxide  (Bi203), 
similar  in  constitution  to  the  acid-forming  arsenious  oxide,  AS2O3, 
possesses  only  basic  properties.  We  will,  therefore,  consider  bismuth 
and  its  derivatives  with  the  metals. 


TABULATION  OF  THE  ELEMENTS  OF  THE  NITROGEN  GROUP. 

The  elements  belonging  here — nitrogen,  phosphorus,  arsenic,  anti- 
mony— present  similar  graded  differences  in  their  physical  and  chemical 
|)r()j)crtics,  just  like  the  elements  of  the  chlorine  and  oxygen  groups,  and 
this  gradation  is  intimately  connected  with  the  atomic  weights.  As 
the  latter  increase  the  substance  condenses,  the  fusibility  and  vola- 
tility decrease,  and  the  similarity  to  the  real  metals  becomes  more  prom- 
inent. 


CARBON  GROUP. 


149 


N 

P 

As 

Sb 

Atomic  weight,  

14.04 

31.0 

75 

120 

Specitic  gravity,  

0.9  (solid) 

1.8-2. 1 

4-7-5-7 

66.7 

Melting  point, 

— 214°  (60 

44" 

red-white  heat 

mm. ) 

Vapor  density  (O^  = 32),  . . 

28.08 

124.0 

150 

240 

These  elements  do  not  resemble  one  another  in  chemical  properties  as 
the  halogens  resemble  one  another,  or  as  sulphur,  selenium  and  tellurium. 
Arsenic  and  antimony  alone  show  chemical  kinship.  What  is  exceed- 
ingly striking  is  the  difference  between  the  remarkable  activity  of  phos- 
phorus in  its  yellow  variety  and  the  sluggish  nitrogen  with  which  red 
phosphorus  may  be  compared.  This  difference  manifests  itself  in  many 
of  their  derivatives  as  has  been  observed  with  their  halides  (NCI3,  NI3 — 
PCI3,  PI3),  and  will  again  be  seen  in  their  oxygen  compounds. 

Excepting  bismuth,  which  is  perfectly  metallic  in  its  nature,  the  ele- 
ments of  this  group  form  gaseous  compounds  with  three  atoms  of  hydrogen. 
Ammonia  (NH3)  possesses  strongly  basic  properties,  and  combines  with 
all  acids  to  yield  ammonium  salts;  phosphine  (PH3)  combines  at  the 
ordinary  temperature  only  with  hydrogen  bromide  and  hydrogen  iodide 
to  form  salt-like  compounds.  Arsine  and  stibine  no  longer  show  basic 
properties.  Arsenic  and  antimony,  as  well  as  the  two  preceding  elements, 
combine  with  the  hydrocarbons  CH3  and  C2H3)  and  form  com- 

pounds which  are  analogous  in  constitution  and  similar  in  character  to 
the  hydrides.  These  compounds  [As(CH3)3  and  Sb(CH3)3]  will  be  de- 
scribed in  Organic  Chemistry ; they  possess  basic  properties  and  yield  salts 
corresponding  to  the  ammonium  salts.  Compounds  of  arsenic  and  anti- 
mony corresponding  to  hydroxylamine  and  hydrazoic  acid  are  not  known. 

The  oxygen  derivatives  of  these  elements  exhibit  a gradation  similar  to 
that  of  the  hydrogen  compounds.  With  increase  of  atomic  weight,  cor- 
responding to  the  addition  of  metallic  character,  the  oxides  which  form 
strong  acids  in  the  lower  series  acquire  a more  basic  nature  in  the  higher 
series. 


CARBON  GROUP. 

The  two  non-metals,  carbon  and  silicon,  and  the  metals,  tin  and  ger- 
manium, comprise  this  group.  They  unite  with  four  atoms  of  hydrogen 
or  with  four  atoms  of  the  halogens. 

1.  CARBON. 

Atom  : C = 12.00. 

Carbon  occurs  free  in  nature  as  the  diamond  and  graphite.  It  consti- 
tutes the  most  important  ingredient  of  all  the  so-called  organic  substances 
originating  from  the  animal  and  vegetable  kingdoms,  and  is  especially 


INORGANIC  CHEMISTRY. 


150 

contained  in  the  fossilized  jiroducts  arising  from  the  slow  decomposition 
of  vegetable  matter — in  i)eat,  in  lignite,  in  bituminous  coal,  and  in  an- 
thracite. In  combination  with  hydrogen  it  forms  the  so-called  mineral 
oils — petroleum  and  asjihaltum.  It  occurs,  further,  as  carbon  dioxide 
(CO^)  in  the  air  and  in  many  waters;  and  in  the  form  of  carbonates 
(marble,  calcite,  dolomite)  comprises  many  minerals  and  entire  rock 
formations. 

It  is  found  in  different  allotropic  modifications  when  free ; these  may 
be  referred  to  the  three  })rincii)al  varieties — diamond,  grajihite  and  amor- 
phous carbon.  In  all  these  forms  it  is  a solid,  even  at  the  highest  tem- 
peratures; non-fusible  and  only  volatile  at  about  3500°  in  the  electric 
arc.  This  deportment  can  only  be  explained  by  the  supi^osition  that  its 
molecules  are  composed  of  a large  number  of  carbon  atoms  (see  j).  106). 
All  the  modifications  of  carbon  are  quite  stable,  but  not  very  reactive. 
When  burned  all  yield  carbon  dioxide. 

1.  The  diamond  occurs  in  alluvial  soils  in  certain  districts  (in  India,  Brazil,  and  South 
Africa)  ; less  frequently  in  itacoluniite,  micaceous  schist,  and  xanthophyllite.  It  was  ob- 
served recently  in  meteoric  iron  from  Canon  Diablo.  It  has  great  luster,  strong  power 
of  refraction,  and  the  greatest  hardness  of  all  substances.  It  crystallizes  in  forms  of  the 
regular  system,  which  are  mostly  rhombic  dodecahedra,  rarely  octahedra.  Ordinarily,  it  is 
perfectly  colorless  and  transparent ; sometimes,  however,  it  is  colored  by  impurities.  Its 
specific  gravity  equals  3.5.  It  does  not  soften  unless  exposed  to  the  most  intense 
heat — between  the  poles  of  a powerful  galvanic  battery.  It  is  then  converted  into  a 
graphitic  mass.  When  heated  in  oxygen  gas  to  700-800°  it  burns  to  carbon  dioxide 
[see  Ber.  23  (1890),  2409].  It  is  scarcely  attacked  at  all  when  acted  upon  by  a mixture 
of  nitric  acid  and  potassium  chlorate. 

The  diamond  has  been  made  artificially.  Molten  iron  dissolves  carbon,  which  sepa- 
rates on  cooling,  mostly  in  the  form  of  graphite.  If  iron,  at  high  temperature,  be 
saturated  with  carbon  and  then  be  quickly  cooled  so  that  its  inner  portions  are  subjected 
to  great  pressure  because  of  the  contraction  of  the  exterior,  upon  sudden  chilling,  the 
carbon  then  crystallizes  in  the  form  of  the  diamond  (Moissan).  Molten  olivine  (mag- 
nesium silicate)  also  dissolves  carbon,  which  again  separates  on  cooling  in  the  form  of 
little  diamond  crystals. 

2.  Graphite  {ypd<pEiv^  to  write)  is  characterized  by  its  oxidation  to  graphitic  acid  when 
it  is  heated  with  a mixture  of  potassium  chlorate  and  nitric  acid.  Like  amorphous 
carbon,  it  is  oxidized  to  mellitic  acid  by  an  alkaline  solution  of  potassium  permanganate, 
or  when  it  is  made  the  positive  electrode  in  the  electrolysis  of  alkaline  solutions.  (See 
Richter’s  Organic  Chemistry.)  Native  graphite  is  found  in  the  oldest  rock  formations, 
and  of  especially  good  quality  at  Altai,  in  Siberia.  It  occurs,  too,  in  considerable  quan- 
tities at  many  places  in  the  United  States.  It  is  occasionally  found  crystallized  in  six-sided 
forms,  but  usually  as  an  amorphous,  grayish-black,  glistening,  soft  mass,  used  in  the 
manufacture  of  lead  pencils.  The  specific  gravity  is  2.25.  It  conducts  heat  and  elec- 
tricity well.  When  away  from  air  contact  it  is  not  altered  even  at  the  highest  tempera- 
tures. It  is  the  only  variety  of  carbon  stable  at  high  temperatures  ; when  heated  the 
other  varieties  pass  into  it.  It  usually  burns  when  heated  in  an  atmosphere  of  oxygen, 
but  with  more  difficulty  than  the  diamond,  forming  carbon  dioxide,  and  leaving  about 
2-5  ])er  cent,  of  ash.  To  purify  the  poorer  and  more  impure  kinds  of  graphite,  they 
are  ])iilverized  and  heated  with  a mixture  of  potassium  chlorate  and  sulphuric  acid  ; the 
{product  is  washed  with  water,  and  the  residue  ignited  (Brody’s  graphite). 

Luzi  claims  two  varieties  of  graphite:  graphiUte  and  the  true  graphite.  The  first 
does  not  swell  up  when  moistened  with  red,  fuming  nitric  acid,  and  is  then  strongly 
ignited  (Bassauer  and  Siberian  graphite).  'I'he  true  graidiites  do  swell  up  under  similar 
tr(“atment  (gra|)hite  from  (feylon,  New  York,  'ficonderoga)  ; see  Ber.  26  ( 1893),  890. 

Urapliile  may  be  obtained  artificially  by  fusing  amoridious  carbon  with  iron  ; when  the 
latter  cools,  a j)ortion  of  the  dissolved  carbon  separates  in  hexagonal  shining  leaflets. 


COMPOUNDS  OF  CARBON  WITH  HYDROGEN. 


15I 

3.  Amorphous  carbon  (a,  privative,  and  /jopcl)//,  form),  is  produced  by  the  carbonization 
of  organic  (containing  carbon)  substances,  anci  is  found  in  a fossilized  state.  Nitric  acid 
and  potassium  chlorate  convert  it  in  the  cold  into  brown  substances  soluble  in  water,  d'he 
purest  amorphous  carbon  is  soot,  which  is  obtained  by  the  imperfect  combustion  of  resins 
and  oils  (like  turpentine)  rich  in  carbon.  Gas  carbon,  called  metallic  carbon,  deposits 
in  the  manufacture  of  gas  in  the  retorts,  and  is  very  hard,  possessing  metallic  luster,  and 
conducting  electricity  well  ; hence  its  use  in  galvanic  batteries.  Coke,  resulting  from  the 
ignition  of  bituminous  coal,  forms  a sintered  mass,  conducting  heat  and  electricity  well. 
Charcoal  is  very  porous,  and  can  absorb  many  gases  and  vapors  ; i volume  of  it  con- 
denses 90  volumes  of  ammonia,  55  volumes  of  hydrogen  sulphide,  and  9 volumes  of 
oxygen.  At  1 00°,  and  under  the  air-pump,  the  absorbed  gases  are  again  liberated. 
Charcoal  will  also  take  up  many  odorous  substances  and  decaying  matter  ; hence  it  is 
employed  as  a disinfectant.  Anhnal  charcoal-  is  obtained  by  the  carbonization  of  animal 
matter  (bones,  blood,  etc.),  and  possesses  the  power  of  removing  many  coloring  sub- 
stances from  their  solutions  ; hence  it  serves  in  the  laboratory  and  in  commerce  for  the 
decolorization  of  dark  solutions. 

All  these  varieties  of  carbon  contain  smaller  or  larger  quantities  of  nitrogen,  hydrogen, 
and  mineral  substances,  which  remain  as  ash  after  combustion.  Hydrochloric  acid  will 
withdraw  almost  all  the  mineral  constituents. 

The  fossil  coal  varieties,  bituminous  coal,  lignite  and  turf,  are  the  products  of  a peculiar, 
slow  decay  of  wood  fiber,  which  gradually  separates  oxygen  and  hydrogen,  and  enriches 
itself  in  carbon.  Fossil  coal  contains  90  per  cent,  and  lignite  70  per  cent,  of  carbon. 
The  fossil  coal  richest  in  carbon,  the  last  product  of  the  alteration,  is  anthracite.  This 
has  lost  all  its  organic  structure,  and  contains  96-98  per  cent,  of  carbon. 


COMPOUNDS  OF  CARBON  WITH  HYDROGEN. 

With  hydrogen,  carbon  forms  an  unlimited  number  of  compounds,  into 
which  all  other  elements,  especially  oxygen  and  nitrogen,  can  enter. 
The  derivatives  of  carbon  have  been  termed  organic  conipoimds,  because 
they  were  formerly  obtained  exclusively  from  vegetable  and  animal 
organisms,  and  the  idea  was  entertained  that  they  were  produced  by  the 
influence  of  forces  other  than  those  forming  the  mineral  substances.  At 
present,  most  carbon  derivatives  are  prepared  artificially  from  the  ele- 
ments by  simple  synthetic  methods;  we  are  aware  that  they  do  not  differ 
essentially  from  mineral  substances.  Hence  the  description  of  the  carbon 
compounds  must  be  arranged  in  the  general  system  of  chemical  bodies. 
This,  however,  is  not  readily  executed  without  sacrificing  the  review  of  a 
defunct  system.  The  derivatives  of  carbon  are  so  numerous,  and  possess 
so  many  peculiarities,  that  it  appears  necessary,  from  a practical  stand- 
point, to  treat  them  apart  from  the  other  compounds,  in  a separate  portion 
of  chemistry,  which  we,  pursuing  the  old  custom,  term  organic  chemis- 
try. We  then  designate  the  chemistry  of  all  other  bodies  as  Inorganic 
Chemistry.  Only  the  simplest  carbon  compounds  will  be  considered  here. 


It  is  only  under  the  influence  of  the  electric  arc  that  the  direct  union 
of  carbon  and  hydrogen  can  be  effected ; the  ])roduct  is  acetylene 
(CJI.J.  All  other  hydrocarbons  are  obtained  indirectly  in  various  ways. 


152 


INORGANIC  CHEMISTRY. 


Methane — ATarsh  Gas — CII^.  "J'his  simplest  liydrocarlion  is  lliat 
cliciiiical  compound  which  is  richest  in  hydrogen  ; it  contains  25.18  jier 
cent.  It  is  lormed  in  the  decay  of  organic  matter  under  water  (in 
swamps,  hence  called  marsh  gas,  and  in  coal  mines),  and  escapes  in 
large  quantities  in  many  regions  of  the  earth  (thus  at  I’.aku,  on  the  Cas- 
pian Sea,  in  Pennsylvania  and  in  Ohio,  where  the  gases  esca])ing  from 
petroleum  wells  contain  as  much  as  90-97  per  cent,  of  methane).  It  is 
produced  in  the  dry  distillation  of  coal,  wood,  i)etroleum  and  other 
organic  substances;  hence  is  a constituent  of  illuminating  gas.  It  may 
be  obtained  synthetically  by  conducting  vapors  of  carbon  bisulphide  and 
hydrogen  sulphide  over  ignited  copper  filings; 

CS2  + 2lI.,S  + 8Cu  4Cu.,S  + CH^. 

For  its  preparation,  heat  a mixture  of  sodium  acetate  with  sodium 
hydroxide : 

CJbOjNa  + NaOH  ^ CII,  -f-  Na,C03. 

Its  formation  from  aluminium  carbide  by  the  action  of  water  is  rather 
interesting.  Aluminium  hydroxide  is  also  produced  (Moissan): 

Al.Cg  + 6II.p  ==  2A1/)3  + 3CIb. 

Methane  is  a colorless,  odorless  gas,  insoluble  in  water.  Its  density 
corres})onds  to  the  molecular  formula  CH^.  Methane,  under  great  pres- 
sure and  at  low  temperatures  (lielow  — 82°,  p.  48),  is  condensed  to  a 
colorless  liquid,  which  boils  at  — 164°  under  the  ordinary  pressure.  Its 
specific  gravity  is  0.415  at  the  boiling  point.  When  ignited  it  burns 
with  a faintly  luminous  flame.  It  affords  a violently  explosive  mixture 
(fire-damp  of  the  miners)  with  two  volumes  of  oxygen  (or  ten  volumes  of 
air) : 

CH^  + 2O2  ^ CO2  -f  2H2O. 

I vol.  2 vols.  I vol.  2 vols. 


MOLECULAR  FORMULA  OF  METHANE.  ATOMIC  WEIGHT  OF  CARBON. 

The  quantitative  analysis  of  methane  shows  that  for  every  i.oi  parts  of  hydrogen  in  it 
there  are  3 parts  of  carbon.  Were  the  formula  CH  (analogous  to  hydrochloric  acid)  then 
the  atomic  weight  of  carbon  would  be  3.  If  it  corresponded  to  the  formula  of  water 


(II20) 

or  ammonia  (NII3)  then  the  atomic  weight  of  carbon  would  equal  6,  or  9 : 

II 

= I.OI  2II  = 2.02 

311  = 

3- 03 

4II  = 4.04 

c 

= 3 C = 6 

C = 

9 

C = 12 

ClI 

= 4.01  CII2  = 8.02 

CII3  = 

12.03 

CIb  = 16.04 

All  these  and  many  other  formulas  agree  with  the  analytical  results.  The  gas  density, 
however,  decides  in  favor  of  the  formula  Clb,  because  a liter  of  methane  weighs  under 
normal  conditions  0.71464  gram  and  this  would  make  the  molecular  weight,  referred  to 
(),  - 32  (.see  p.  79),  = .6.0. 

d hat  the  atomic  weight  of  carbon  is  really  12,  is  proven  by  the  fact  that  of  all  its 
in nunicrahle  derivatives,  not  one  contains  in  its  molecule  less  than  12  farts,  by  7veight 
{^therefore  in  the  mol-volume  (p.  98)  not  less  than  12  grams'),  of  this  element.  It  follows, 
with  certainty,  from  the  i)criodic  system  of  elements. 


ETHANE ETHYLENE.  I53 

From  the  formula  CII^  it  follows  that  in  i volume  of  methane  there  are  2 volumes  of 
hydrogen : 

CH^  contains  2li.^ 

I vol.  2 vols 

This  is  proved  indirectly  by  the  combustion  of  methane  with  oxygen  in  a eudiometer 
(see  p.  137)-  Four  atoms  of  hydrogen  yield  two  molecules  of  water  ; one  atom  of  carbon 
yields  one  molecule  of  carbon  dioxide,  CO2  Hence  the  volume  relation  in  the  combus- 
tion of  methane  in  oxygen  is  expressed  by  the  equation ; 

CH,  + 2O2  =•  CO2  + 21120. 

I vol.  2 vols.  1 vol.  2 vols. 

In  2 volumes  of  aqueous  vapor  there  are  2 volumes  of  hydrogen  ; hence  in  I volume 
of  Cll^,  there  are  2 volumes  of  II2.  The  result  of  the  eudiometric  analysis  conhrms 
these  conclusions 


Ethane  or  Ethyl  Hydride,  C2Hg,  is  formed  when  hydrogen  m 
statu  nascendi  acts  upon  ethyl  bromide : 

C2H5Br  + H2  =■  + HBr, 

or  by  the  action  of  potassium  or  sodium  upon  methyl  iodide  (p.  159) : 

2CH3I  -f  Na2  = C2II6  + 2NaI. 

This  is  a colorless  gas,  insoluble  in  water,  and  when  ignited  it  burns 
with  a feebly  luminous  flame.  Liquid  ethane  boils  at  — 89.5°  under  735 
mm.  pressure. 

Besides  methane  (Cbb)  and  ethane  (C2Hg)  there  exists  a long  series  of  hydrocarbons 
of  the  general  formula  CnH2n-l-2  g’,  Cgllg,  C5IIJ2.  etc.),  in  which  each  mem- 

ber differs  from  the  preceding  and  next  following  by  iC  and  2H  (CH2).  Bodies 
belonging  to  such  a series,  greatly  alike  in  their  chemical  behavior,  are  termed  honiolognes 
{ouoiog,  similar;  Aoyo^,  nature  of  the  affair).  In  addition  to  this  series  of  satm-ated 
hydrocarbons  others  exist,  with  less  hydrogen,  and  by  the  addition  of  the  latter,  pass 
into  the  saturated^  and  may,  therefore,  be  termed  jmsaturated.  The  first  uusattirated 
.series  has  the  formula  Cnll^,,,  the  second,  the  fonnula  CnH2n— 2,  etc.  The  lowest  member 
of  the  series  Cnll2n  is  ethylene  (see  Chemical  Structure,  p,  168). 

Ethylene,  C2H^,  is  formed  in  the  destructive  distillation  of  wood, 
bituminous  coal,  and  many  carbon  compounds,  hence  is  contained  in 
illuminating  gas.  It  is  most  easily  obtained  by  the  action  of  sulphuric 
acid  upon  alcohol,  whereby  the  acid  withdraws  the  elements  of  water  from 
the  latter  • 

C2lTgO  — 1120  =r  C2IL. 

Alcohol.  Ethylene. 

It  is  a colorless  gas,  of  weak,  ethereal  odor.  Its  critical  temperature 
is  its  critical  pressure,  51  atmospheres.  Under  a pres- 

sure of  I atmosphere,  licpiid  ethylene  boils  at  — 102°,  and  when  in  a 
vacuum  at  ■ — 150°.  It  solidifies  at  — t6o°.  As  it  does  not  solidify  on 
va])orizing,  it  is  well  adapted  for  the  liciuefaction  of  other  gases  (p.  48). 
It  burns  with  a bright,  luminous  flame,  decomposing  first  into  marsh  gas 
and  free  carbon  : 

€211,  = CII,  + C. 

The  methane  then  burns  and  heats  the  i)articles  of  carbon  in  the  flame 


154 


INORGANIC  CHEMISTRY. 


to  incandescence  ; tliese  are  then  consumed  to  carl)on  dioxide  (CO.^)  (see 
l>er.  26  ( I ^93),  III,  9). 

'riie  unsaliirated  compound,  ethylene,  unites  directly  with  two  atoms 
of  chlorine  and  of  bromine: 

Cdb  b Cl,  C,TI,C1,. 

The  resulting  compounds,  and  C,l  l^Ilr.^,  are  oily  liciuids;  hence 

the  name  olefiant  ^as  for  ethylene. 

The  first  member  of  the  second  unsaturated  series  is  acetylene, 

Acetylene  is  juoduccd  in  the  dry  distillation  of  many  carbon  com- 
pounds, and  is  present  in  coal  gas,  to  wh.ch  it  imparls  a ])eculiar  ]>ene- 
trating  odor.  It  is  also  formed  in  the  incom])lete  combustion  of  coal  gas — 
e.  when  the  fiame  of  a Bunsen  burner  strikes  back.  Berthelot  syn- 
thesized acetylene  by  causing  the  electric  spark  to  strike  across  from  car- 
bon points  in  an  atmosphere  of  hydrogen  : 

C,  + II,  = C,H,. 

No  other  hydrocarbon  has  as  yet  been  directly  built  up  from  its  ele- 
ments. 

The  production  of  acetylene  by  the  action  of  water  or  dilute  acids 
upon  metallic  carbides  is  very  interesting,  e.  g.,  from  calcium  carbide: 

CaC,  + 211,0  = CjII,  + Ca(OH),. 

Pure  acetylene  is  a colorless  gas  with  a peculiar  odor.  It  can  be  readily 
condensed.  Its  critical  temperature  is  37°;  its  critical  pressure  equals 
68  atmospheres.  At  0°  it  may  be  liquefied  by  a pressure  of  21.5  atmos- 
pheres. Liquid  acetylene  boils  at  — 83°.  When  it  is  poured  it  changes  in 
part  to  a snow-like,  combustible  mass.  At  19.5°,  i kilogram  of  liquid  acet- 
ylene will  yield  896  liters  of  gas  under  the  ordinary  atmospheric  pressure. 
At  medium  temperatures  it  dissolves  in  an  equal  volume  of  water;  there- 
fore, it  must  be  collected  over  strong  salt  solutions.  It  is  very  soluble  in 
acetone,  i volume  of  the  latter  dissolving  25  volumes  of  acetylene  at  the 
ordinary  temperature  and  pressure. 

Being  an  unsaturated  hydrocarbon  it  combines  with  chlorine;  with 
two  or  four  atoms  of  the  same : 

C2H2CI2,  C2H2Ch. 

It  is  chemically  very  much  like  hydrazoic  acid  (p.  132).  It  possesses  acid 
l)roperties  but  these  are  not  so  pronounced  as  in  the  case  of  hydrazoic 
acid ; acetylene  and  many  of  its  metallic  derivatives  are  explosive  (p.  31). 

As  a result  of  o])servations  made  by  Willson,  an  American,  when 
attempting  to  ])repare  the  alkaline  earth  metals  by  the  reduction  of  their 
oxides  with  carbon  in  the  electric  furnace,  calcium  carbide  has  been  made 
since  1894  on  a large  scale,  for  the  purpose  of  producing  acetylene  by  the 
method  employed  by  Moissan.  As  acetylene  is  now  so  readily  accessible 
it  plays  an  imj)ortant  role  as  an  illuminant.  It  burns,  on  issuing  from  a 
small  aperture  under  a definite  i)ressure,  with  a blinding  white  flame, 
almost  free  from  soot.  As  it  is  readily  condensed  it  is  especially  adapted 
for  the  illumination  of  vessels,  trains,  etc. 


THE  NATURE  OF  FLAME. 


155 


Mixtures  of  acetylene  with  from  1.25  to  20  volumes  of  air  are  explosive; 
a most  powerful  explosion  will  occur  when  the  j^roportion  is  i volume  of 
acetylene  and  12  volumes  of  air.  The  acetylene  prejjared  from  commer- 
cial calcium  carbide  usually  contains  hydrogen  sulphide  and  phosphide. 

Compare:  Fr.  Liebetanz,  Calciumkarbid  und  Acetylentechnik,  2. 
Aufl.,  Leipzig,  1899;  F.  Dommer,  Calciumkarbid  und  Acetylen ; 
deutsch  von  Landgraf,  Miinchen  und  Leipzig,  1898. 

The  three  hydrocarbons  considered  above,  methane  (CHJ,  ethylene 
(C2HJ,  and  in  slight  amount  acetylene  (C2H2),  constitute,  together  with 
hydrogen  and  carbon  monoxide  (CO),  ordinary  ilhwiinating gas,  which  is 
produced  in  the  dry  distillation  of  bituminous  coal,  lignite,  or  wood. 
The  illuminating  power  is  influenced  by  its  quantity  of  ethylene  and  acet- 
ylene (and  their  homologues). 

Late  investigations  indicate  that  all  hydrocarbons  are  broken  down  by  the  heat  of  the 
flame  into  acetylene  and  carbon,  which  then  burn  and  become  luminous  (see  Z.  f.  anorg. 
Ch.  9 (1895),  233). 


THE  NATURE  OF  FLAME. 

We  are  aware  that  every  chemical  union  which  occurs  in  a gaseous 
medium,  and  is  accompanied  by  the  evolution  of  light  is  designated  com- 
bustion. Some  bodies,  like  sulphur  and  phosphorus,  yield  a flame  when 
burned  in  the  air  or  in  other  gas;  such  sub- 
stances are  converted  into  gases  or  vapors  at  the 
temperature  of  combustion.  Pure  carbon  burns 
without  a flame,  becomes  incandescent,  because 
it  is  non-volatile.  The  carbon  compounds, 
wood,  bituminous  coal,  and  tallow,  are,  indeed, 
not  volatile,  but  burn  with  a flame  because 
under  the  influence  of  heat  they  develop  com- 
bustible gases.  Flame  is,  therefore,  nothing 
more  than  a combustible  gas  heated  to  incan- 
descence. We  know,  too,  that  hydrogen  burns 
in  oxygen  and  chlorine,  conversely,  oxygen  and 
chlorine  will  burn  in  hydrogen  (p.  57),  and 
that  illuminating  gas  burns  in  the  air,  therefore 
air  (its  oxygen)  burns  in  the  former.  This  may 
be  demonstrated  in  the  same  manner  as  in  the 
case  of  chlorine  and  hydrogen. 

The  relative  combustibility  and  the  so-called  return  of 
the  flame  may  V)e  very  plainly  illustrated  by  means  of  the 
following  contrivance  : An  ordinary  lamp  chimney  (Fig. 

57)  is  closed  at  its  lower  end  with  a cork,  through  which  FlG.  57. 

two  tubes  enter ; the  narrow  lube,  a,  somewhat  con- 
tracted at  its  end,  is  connected  with  a gas  stop-cock  ; the  other  tube,  h (best  a cork-borcr \ 
is  about  5 mrn.  wide,  and  communicates  with  the  air.  The  gas  issuing  from  the  tube  a 
is  ignited,  and  the  chimney  is  then  drojiped  over  the  not  too  large  flame  ; it  contituies  to 
burn  along  (|uietly,  as  sufficient  air  enters  through  the  wide  tube  b.  Upon  increasing  the 
supply  of  gas,  the  flame  becomes  larger,  the  globe  fills  with  illuminating  gas,  while  the 


156 


INORGANIC  CHEMISTRY. 


air  is  displaced.  The  p;as  flame  is  extin^uislied,  and  an  air  flame  appears  upon  the 
wider  tube,  h,  as  the  enlerinjj;  air  continues  to  burn,  in  the  atmosphere  of  illuminatinfr 
f^as.  'I'he  exce.ss  of  tlie  latter  e-scapinj^  from  the  upper  ])ortion  of  the  ^lobe  may  be 
if^nited,  and  we  then  have  a gas  flame  above,  while  within  the  globe  we  have  an  air  llame. 
On  again  lessening  the  gas  How  the  air  flame  will  distribute  itself,  extend  to  the  exit  of 
the  tube  and  then  the  gas  flame  will  appear  upon  the  latter,  while  the  flame  above  the 
globe  is  extinguished.  In  this  manner,  we  may  repeat  the  return  ])roeess  of  flames  at 
will.  That  the  air  actually  burns  in  the  air  flame  may  be  plainly  proved  if  we  introduce 
a small  gas  flame  from  c,  through  the  wide  metallic  tube the  little  flame  will  con- 
tinue to  burn  in  the  air  flame,  but  will  be  extingui.shed  if  it  be  introduced  higher  up 
into  the  atmosphere  of  illuminating  gas. 

We  say  ordinarily  that  only  those  bodies  are  combustible  which  burn 
in  an  atmosphere  of  oxygen  or  in  air.  If  we  imagine,  however,  an  at- 
mos])here  of  hydrogen,  or  illuminating  gas,  then  bodies  rich  in  oxygen 


Fig.  58. 


must  be  combustible  in  these.  In  fact,  nitrates,  chlorates,  etc.,  burn  in 
an  atmosphere  of  illuminating  gas  with  the  production  of  an  oxygen 
flame.  This  may  be  demonstrated  as  follows:  An  Argand  lamp  chimney 
(Fig.  58)  is  closed  at  its  lower  end  with  a cork,  bearing  a gas-conducting 
tube,  d'he  gas  which  escapes  through  the  opening  of  the  metal  cover, 
a,  is  ignited.  Then  the  substance  (potassium  or  barium  chlorate,  etc.) 
is  introduced  into  the  flame  on  an  iron  spoon  provided  with  a long 
handle,  heated  to  the  temperature  of  decomposition  (disengagement  ot 
oxygen),  and  the  sjioon  then  jilunged  through  the  opening  into  the  gas 
atmosjihere.  d'he  substance  burns  with  a brilliant  light,  and  gives  a 
cliaracteristic  flame  reaction. 


THE  DENSITY  OF  THE  GAS  OF  FLAMES. 


157 


The  brilliancy  or  luminosity  of  a flame  is  influenced  by  the  nature  of 
the  substances  contained  in  it,  also  by  their  temperature  and  density.  In- 
candescent gases  shine  very  faintly  per  se ; this  is  especially  true  when 
they  are  diluted.  Thus  hydrogen,  ammonia  and  methane  burn  with  a pale 
flame.  Even  sulphur  burns  in  the  air  with  a slightly  luminous  flame.  If,  on 
the  contrary,  sulphur  or  phosphine  be  permitted  to  burn  in  pure  oxygen, 
or  arsenic  and  antimony  in  chlorine  gas,  an  intense  display  of  light  fol- 
lows. This  depends  on  the  fact  that  the  flame  is  not  diluted  by  the  nitro- 
gen of  the  air,  and  therefore  develops  a higher  temperature.  That  the 
density  of  the  gas  of  flames  exercises  a great  influence  upon  the  lumi- 
nosity is  proved  by  the  fact  that  hydrogen,  compressed  into  a smaller 
space  with  oxygen,  burns  with  intense  light. 

A slightly  luminous  flame  may  be  rendered  intense  by  introducing  solid 
particles  into  it.  For  example,  if  hydrogen  be  passed  through  liquid 
chromium  oxychloride  (Cr02Cl2)  it  burns  with  a bright,  lumi- 
nous flame,  because  the  chromium  oxychloride  is  changed  by 
the  high  temperature,  with  the  absorption  of  hydrogen,  into 
water,  hydrochloric  acid,  and  solid,  non-volatile  chromium 
oxide,  whose  particles  are  heated  to  incandescence  by  the 
hydrogen  flame.  The  illuminating  power  of  the  various 
hydrocarbons  and  carbon  compounds  is  similarly  explained. 

Marsh  gas,  CH^,  and  ethane,  C2Hg,  give  a pale  flame,  be- 
cause they  burn  directly  to  aqueous  vapor  and  carbon  dioxide ; 
ethylene,  on  the  contrary,  burns  with  a bright,  luminous 
flame,  because,  by  the  temperature  of  combustion,  it  decom- 
poses first  into  methane  and  carbon,  whose  particles  glow  in 
the  flame  (see  p.  153). 

Let  us  consider  the  flame  of  an  ordinary  stearin  candle : On 
approaching  the  wick  with  a flame  the  stearin  melts,  is  drawn 
up  by  the  fibers  and  converted  into  gaseous  hydrocarbons, 
which  ignite,  and  by  their  chemical  union  with  the  oxygen 
of  the  air  produce  the  flame.  Three  zones  are  distinguish- 
able in  this  flame.  The  unaltered  gases  exist  in  the  inner  non-volatile 
zone  a (Fig,  59);  they  cannot  burn  because  of  lack  of  air.  If  the 
lower  end  of  a thin  glass  tube  be  inserted  here  the  gases  will  rise  in  it, 
and  may  be  ignited  at  the  upper  end.  There  is  a partial  combustion  of 
the  gases  in  the  middle,  luminous  part,  y,  g;  ethylene,  C2H4,  breaks 
down  here  into  methane,  CH^,  and  carbon,  C ; the  first  burns  completely, 
while  the  carbon  is  heated  to  a white  heat,  because  there  is  not  sufficient 
oxygen  ])resent  for  its  combustion.  The  presence  of  carbon  particles  in 
the  luminous  part  may  be  easily  proved  by  placing  a cold  glass  rod  or  a 
wire  in  it ; it  will  at  once  be  covered  with  soot. 

In  the  outer,  very  feebly  luminous  and  almost  invisible  mantle,  bj  c,  d, 
of  the  flame,  which  is  completely  surrounded  by  air,  occurs  the  perfect 
combustion  of  all  the  carbon  to  carbon  dioxide. 

A perfectly  identical  structure  is  possessed  by  the  ordinary  illuminating 
gas  flame.  By  bringing  as  much  air  or  oxygen  into  it  as  is  necessary  for 
the  complete  combustion  of  all  the  carbon,  none  of  the  latter  separates 
(see  p.  158),  and  there  is  produced  a faintly  luminous  but  very  hot  flame. 


Fig.  59. 


INORGANIC  CHEMISTRY. 


158 


Upon  this  ])rincii)le  is  based  tlie  construction  of  the  Bunsen  burner,  the 
flame  of  which  is  employed  in  laboratories  for  heating  and  ignition. 
Fig.  60  represents  a form  of  the  same.  The  upi)er  tube,  c,  is  screwed 
into  the  lower  portion,  and  in  the  figure  is  separated  merely  for  the  sake 
of  exi)lanation.  The  gas  enters  through  the  narrow  ojiening,  a,  from  the 
side  gas  tube,  and  mixes  with  air  in  the  tube  c,  which  enters  through 
the  openings  of  the  ring,  b.  In  this  way  we  obtain  a flame  which  is  but 
faintly  luminous,  although  affording  an  intense  heat.  On  closing  the 
openings  in  b the  air  is  cut  off,  and  the  gas  burns  at  the  uiiper  end  of  the 
lube  c with  a bright,  strongly  smoking  flame.  The  non-luminous  flame 
contains  an  excess  of  oxygen,  and  hence  oxidizes — oxidizing It 
is  employed  to  effect  oxidation  reactions.  Its  temperature  is  about  i 200°. 
'I'he  luminous  flame,  on  the  other  hand,  is  reducing  in  its  action  (tem- 
])erature  1000°  C.),  and  is  designated  the  7'educing  flame,  because  the 
glowing  carbon  in  it  abstracts  oxygen  from  many  substances. 


The  non-lurninosity  of  tlie  Bunsen  burner  flame,  clue  to  addition  of  air,  depends  on  a 
more  complete  combustion  of  the  separated  carbon  or  of  the  yet  undecomposed  hydro- 
carbons. 

Another  variety  of  non-luminosity  of  hydrocarbon  flames  is  induced  by  the  ad- 
mixture of  inactive  gases,  like  nitrogen  and  carbon  dioxide.  By  this  means  the 

flame  is  enlarged  and  the  combustion,  as  in  the  luminous  flame,  takes  place  only 
in  the  outer  cone  ; further,  the  temperature  is  lowered,  and  probably  does  not 
acquire  the  decomposition  temperature  of  methane  and 
ethylene.  The  simple  extension  of  an  illuminating  flame 
upon  a plate,  will  render  it  non-luminous,  because  then 
the  air  comes  in  contact  with  a larger  flame  surface.  On 
heating  a gas  made  non-luminous  by  the  admixture  of 
nitrogen,  and  then  letting  it  burn,  its  flame  becomes  lumi- 
nous because  the  increased  temperature  can  induce  the 
decomposition  of  methane. 


rh-.  : 


irisfai 


Fig.  60. 


Fig.  61. 


In  rendering  flame  non-luminous  by  carbon  dioxide,  we  must  also  consider  that  the 
same  is  converted,  by  the  particles  of  carbon,  into  carbon  monoxide  : 

CO2  -f  C = 2CO. 

Indeed,  but  a few  per  cent,  of  carbon  dioxide  in  a gas  flame  suffices  to  considerably 
diminish  its  luminosity  : 

(',H,  I ro.,  _ C’ll,  + 2C(), 

1 vul.  I vul.  I vul.  2 vols. 
while  the  presence  of  nitrogen  is  far  less  detrimental. 


COMPOUNDS  OF  CARBON  WITH  THE  HALOGENS.  1 59 

Every  substance  requires  a definite  temperature  for  its  ignition — tem- 
perature of  igjiition.  When  a substance  is  once  ignited  it  generally  burns 
further,  because  additional  particles  are  raised  to  the  temperature  of 
ignition  by  the  heat  of  combustion.  By  rapid  cooling  (<?.  g.,  by  the 
introduction  of  a piece  of  metal  into  a small  flame)  every  flame  may  be 
extinguished.  By  holding  a wire  gauze  over  the  opening  of  a gas  lamp, 
from  which  gas  issues,  and  igniting  the  same  above  the  wire  (Fig.  6i), 
the  latter,  being  a good  conductor  of  heat,  cools  the  flame  so  much  that 
it  is  incapable  of  igniting  the  gas  below  the  gauze.  Upon  this  phe- 
nomenon depends  the  construction  of  Davy’s  safety  lamp,  which  is  used 
in  coal  mines  to  avoid  ignition  of  the  fire-damp.  It  is  an  ordinary  oil 
lamp  surrounded  and  shut  off  from  the  air  by  a metallic  wire  gauze.  On 
bringing  a lighted  lamp  of  this  sort  into  an  explosive  mixture,  or  into  a 
combustible  gas  {e.  g.,  into  a large  jar,  in  which  ether  is  present),  the 
gas  penetrating  into  the  lamp  will  burn,  but  the  combustion  will  not  ex- 
tend to  the  external  gases. 


COMPOUNDS  OF  CARBON  WITH  THE  HALOGENS. 

Carbon  combines  directly  with  fluorine  alone  of  the  halogens,  forming 
gaseous  carbon  tetrafluoride,  CFl^.  Compounds  with  all  the  halogens 
result,  however,  by  the  action  of  the  latter  upon  the  hydrocarbons.  We 
have  seen  that  chlorine  and  bromine  act  upon  water,  ammonia,  hydrogen 
sulphide,  etc.,  in  such  manner  as  to  unite  with  the  hydrogen  to  form 
hydrogen  chloride,  etc.,  while  the  other  element  is  either  set  free  or  is 
also  combined  with  the  chlorine  (p.  128).  Chlorine  acts  similarly  upon 
the  hydrocarbons;  here  hydrogen  is  displaced,  atom  after  atom,  by 
chlorine,  forming  hydrogen  chloride  and  chlorine  derivatives : 

CH,  + CI2  = CH3CI  -f  HCl, 
etc  + 2C1,  = ciccq  + 2HC1, 

CH^  -p  3CI2  =:  CHCI3  -f  3HCI,  etc. 

Such  a process  is  termed  substitution^  and  the  products  substitution  prod- 
ucts. These  halogen  derivatives  are  distinguished  from  the  halides  of 
other  non-metals  by  their  comparatively  greater  stability  toward  water 
(see  pp.  148,  1 61).  In  this  way  we  obtain  from  methane,  CH^,  the 
products,  CH3CI,  CH2CI2,  CHCI3  (chloroform)  and  finally  CCl^,  carbon 
tetrachloride,  a colorless,  ethereal  liquid,  boiling  at  76°. 

The  compound  hexachlorethane,  obtained  by  the  action  of  chlorine  upon  ethane, 

CjHg,  is  a crystalline  mass,  melting  and  boiling  at  186°.  On  conducting  its  vapors  through 
a red-hot  tube  they  decompose  into  ethylene  tetrachloride  and  chlorine : 

QCl,  = C,Ch  -p  cp. 

Ethylene  tetrachloride^  CjCp,  is  a liquid  boiling  at  122°. 

P'luorine,  bromine,  and  iodine  yield  similar  compounds  ; they  will  be  treated  more 
exhaustively  in  Organic  Chemistry. 


i6o 


INORGANIC  CHRMISTRY. 


Tlie  heals  of  formation  of  (lie  previously  mentioned  hydrocarbons  (from  amorphous 
carbon  and  hydrogen)  are  deduced  from  tlie  heat  of  conil)ustif)n,  and  ef|ual  : 

(C,II,)  = 2I.7;  (C„ig  = 28.6;  (C„nj  --2.7;  (C2, II,) -47-8. 

The  a])sorption  of  heat  in  the  formation  of  acetylene,  f',!!,,  and  ethylene,  is 

explained  by  the  fact  that  the  solid  carbon  molecules  must  first  be  gasi(ie<I  and  sejiarated 
into  atoms  in  order  to  unite  with  hydrogen.  The  energy  reijuisite  for  this  eipials  very 
nearly  39.7  Cal.  for  12  parts  by  weight  of  amorjihous  carbon  (76.4  Cal.  for  24  jiarts)  ; 
then  the  preceding  heats  of  formation  must  be  increased  about'this  amount,  in  order  to 
express  the  true  heat  of  formation  (from  gaseous  carbon  atoms)  of  the  hydrocarbons. 

Beginning  with  acetylene,  C,!!,,  we  discover  that  its  conversion  into  ethylene,  C,!!^, 
and  into  ethane,  C,IIg,  ensues  with  heat  disengagement : 

(c,ii„h.,)  = 45-i;  ((’2^4,11,)  = 25.9 

In  accordance  witli  this  acetylene  readily  unites  with  liydrogen  (/w  statu  naseendi,  or 
by  action  of  platinum  sponge),  forming  ethylene  and  elhane. 

The  heats  of  formation  of  the  carbon  chlorides  approach  those  of  the  hydrogen  deriva- 
tives very  closely  ; 

= 21.0;  (C2,Cl^-gas)  = —I. I 


2.  SILICON. 

Atom  : Si  = 28.4. 

Silicon  next  to  oxygen  is  the  most  widely  distributed  element  in  nature. 
Owing  to  its  affinity  for  the  latter  it  does  not  occur  in  a free  condition. 
Combined  with  oxygen  as  silicon  dioxide  (Si02),and  in  the  form  of  salts 
of  silicic  acid  (silicates)  it  comprises  many  minerals  and  almost  all  the 
crystalline  rocks. 

It  was  first  obtained  in  a free  condition  by  Berzelius  (1823)  on  heating 
potassium  silico-fluoride  (K2SiFlg)  with  metallic  potassium: 

K,SiFle  + 4K  = 6KF1  + Si. 

The  ignited  mass  is  treated  with  water,  which  dissolves  the  potassium 
fluoride  and  leaves  the  silicon  as  a brown,  non-lustrous,  amorphous  pow- 
der. It  can  also  be  obtained  by  heating  silicon  fluoride  (SiFl^)  with 
sodium,  or  very  readily  by  igniting  a mixture  of  quartz  sand  (silicon  di- 
oxide) with  magnesium  powder  [see  Cl.  Winkler,  Ber.  23  (1890),  2652]: 

2SiO.,  -f  2Mg  = Si  -f  2MgO. 

It  burns  to  silicon  dioxide  (SiOg)  when  heated  in  the  air. 

Another  modification — the  crystalline  silicon — discovered  by  St.  Claire 
Deville,  is  obtained  by  fusing  sodium  silico-fluoride  with  aluminium  or 
with  sodium  and  zinc.  The  separated  silicon  dissolves  in  the  molten 
zinc,  and  on  cooling,  deposits  out  in  crystals,  which  remain  on  dissolving 
the  zinc  in  hydrochloric  acid.  The  crystalline  product  consists  of  black, 
shining  octahedra  and  needles,  of  specific  gravity  2.49,  and  of  very  great 
hardness.  U])on  ignition  in  the  air  or  oxygen  it  is  not  oxidized;  it  is  not 


SILICON. 


l6l 


attacked  by  acids.  On  boiling  it  with  a sodium  or  potassium  liydroxide 
solution  it  dissolves,  forming  a silicate  and  liberating  hydrogen  : 

Si  + 4KOH  = K,SiO,  + 211,. 

Heated  in  chlorine  gas,  silicon  burns  to  the  chloride. 

Hydrogen  Silicide,  SiH^,  the  analogue  of  CH^,  is  produced  like 
arsine  and  stibine  by  dissolving  an  alloy  of  silicon  and  magnesium  in 
dilute  hydrochloric  acid : 

Mg, Si  + 4HCI  = SiH^  + 2MgCl2. 

Such  an  alloy  can  be  obtained  by  heating  finely  pulverized  quartz  or 
glass  with  1^4  times  its  quantity  of  magnesium  powder.  The  escaping 


hydride  contains  admixed  hydrogen,  has  a disagreeable  odor,  ignites 
spontaneously  in  the  air,  and  burns  to  the  dioxide  and  water : 

SiH^  + 20,  = SiO,  + 2H2O. 

Perfectly  pure  silicide,  free  from  hydrogen,  is  obtained  by  heating  the 
compound,  SiH(O.C2H5)3,  which  will  be  described  in  Organic  Chemistry. 
In  the  air  at  ordinary  pressure  it  ignites  only  upon  warming;  if,  however, 
the  gas,  by  diminution  of  pressure  or  by  the  addition  of  hydrogen,  is 
diluted,  it  becomes  spontaneously  combustible  at  the  ordinary  tempera- 
ture. A red  heat  decomposes  the  hydride  into  amorphous  silicon  and 
hydrogen.  When  mixed  with  chlorine  it  inflames  and  probably  forms 
substitution  products  similar  to  those  of  methane  (CHJ.  Pure  hydrogen 
silicide  condenses  to  a liquid  at  — 1°  and  a pressure  of  100  atmospheres. 

Silicon  Chloride,  SiCl^,  results  from  the  action  of  chlorine  upon 
silicon,  or  magnesium  silicide,  or  by  conducting  chlorine  over  an  ignited 
mixture  of  the  dioxide  and  carbon  (Fig.  62) : 

SiO,  + 2C  -f-  2CI,  = SiCh  + 2CO. 


14 


i62 


MORGAN  1C  CHEMISTRY. 


The  niixtiirc  is  i)laccd  in  a porcelain  tube,  wliicli  is  licated  lo  a red 
heat.  Tlie  chlorine  generated  in  the  flask  is  dried  in  the  two  calcium 
chloride  towers  and  the  silicon  tetrachloride  is  condensed  in  the  receiver 
chilled  with  ice-water.  While  carbon  or  chlorine  does  notact  separately 
upon  the  dioxide,  when  they  act  simultaneously  the  reaction  is  induced  by 
the  mutually  supporting  affinities  of  carbon  for  oxygen  and  of  chlorine 
for  silicon. 

Silicon  chloride  is  a colorless  liquid,  having  a sj^ecific  gravity  of  1.5, 
and  boiling  at  59°.  It  fumes  in  the  air,  and  is  decomposed  by  water 
into  silicic  and  hydrochloric  acids: 

SiCh  + 4lbO  = IbSiO,  q-  4IICI. 

This  is  a remarkable  distinction  between  it  and  carbon  tetrachloride. 


Silicon  chloride  may  he  employed  in  determining  the  atomic  weight  of  silicon.  Analysis 
shows  that  it  contains  7.1  parts  of  silicon  for  every  35.45  parts  of  chlorine.  .Siij)posing, 
from  the  great  analogy  of  the  silicon  compounds  to  tho.se  of  carbon,  that  its  formula  is  SiCfi, 
the  atomic  weight  of  the  silicon  would  be  28.4  : 

Si  = 28.4 

Cl,  = 141.80  (4  X 35-45)- 
SiCl,  — 170.20 

This  .supposition  is  confirmed  by  the  vapor  density  of  the  chloride  and  other  deriva- 
tives of  silicon. 


Silicon  Bromide,  SiBr,,  and  Silicon  Iodide,  Sil„  are  formed 
in  the  same  manner  as  the  chloride.  The  first  is  a colorless  liquid,  of 
s])ecific  gravity  2.8,  becoming  solid  at  — 12°  and  boiling  at  +153°. 
The  iodide  forms  colorless  octahedra,  melting  at  120°  and  boiling  at  290°. 
Like  the  chloride,  both  are  decomposed  by  water. 

Besides  these  compounds,  which  may  be  viewed  as  hydrogen  silicide,  in  which  all  of  the 
hydrogen  is  replaced  by  halogens  (see  p.  159),  others  exist  in  which  only  a part  of  this 
element  is  replaced.  Thus,  silicon  chloroform ^ SiHClg,  corresponds  to  the  chloroform 
(CHClg)  derived  from  methane.  It  is  produced  by  the  action  of  phosphorus  pentachlo- 
ride,  or  antimony  pentachloride,  on  hydrogen  silicide : 

SiH,  + 3SbCl5  = SiHCb  + 3SbCl3  + 3HCI ; 

or  upon  heating  silicon  in  dry  hydrogen  chloride  gas  : in  this  case  a mixture  of  the  tetra- 
chloride and  silicon  chloroform  results  (.see  also  Ber.  22,  188).  These  compounds  may 
be  separated  by  fractional  distillation.  Silicon  chloroform  is  a colorless  liquid,  of  specific 
gravity  1.6,  and  boils  at  35-37°.  The  vapor  density  corre.sponds  to  the  molecular 
formula  .SillClj.  It  fumes  in  the  air,  and  decomposes  with  water  into  silicic  and  hydro- 
chloric acids. 

The  silicon  bromoform,  SillBrj,  and  iodoform,  SiHIj,  are  very  similar  to  chloroform  ; 
they  corre.spond  to  the  analogous  carbon  compounds. 

The  compounds  Si^Clg  and  Si.^Ig  are  known.  They  correspond  to  hexachlorethane, 
CqClg.  From  all  the.se  data  we  ob.serve  the  analogy,  based  on  formulas  alone,  between 
.silicon  and  carbon. 


Silicon  Fluoride,  SiFl„ 
silica  : 


SiO.,  ^ 


is  formed  when  hydrofluoric  acid  acts  upon 


4IIFI  :=  SiFl,  -P  211,0. 


To  y)reparc  it,  a mixture  of  fluorite  and  ])owdered  glass,  or  sand  (SiO.,), 
is  warmed  with  suliduiric  acid  ; by  the  action  of  sulphuric  acid  upon  the 


SILICON. 


163 


fluorite  h3'drogen  fluoride  (p.  64)  is  disengaged,  and  this  reacts  upon  the 
silicon  dioxide  in  tlie  manner  indicated  in  the  above  equation.  I'he  liber- 
ated gas  is  collected  over  mercury. 

It  is  colorless,  has  a penetrating 
odor,  and  fumes  strongly  in  the 
air.  Its  vapor  density  corresponds 
to  the  molecular  formula  SiFl^. 

Its  deportment  with  water  is  very 
characteristic ; it  is  decomposed 
thereby  into  silicic  acid  (H^SiOj 
and  hydrogen  silico-fluoride : 
sSiFl,  + 4H,0  = H.SiO,  + 

2H2SiFlg.  For  the  execution  of 
this  method,  conduct  the  tetra- 
fluoride  through  a glass  tube 
into  a vessel  containing  water 
(Fig.  63).  Gelatinous  silicic  acid 
separates  out,  and  the  gaseous 
hydrogen  silico-fluoride  (H^SiFlg) 
remains  dissolved  in  the  water. 

As  the  separating  silicic  acid 
may  easily  obstruct  the  opening 
of  the  glass  tube,  the  latter  is 
allowed  to  project  a slight  distance  into  mercury.  The  solid  silicic  acid 
is  separated  from  the  aqueous  solution  by  filtration. 

Hydrogen  Silico-fluoride,  H2SiFlg,  is  only  known  in  aqueous  solu- 
tion. Upon  evaporation  at  a low  heat  it  decomposes  into  the  tetrafluor- 
ide  and  hydrofluoric  acid(SiFl^  and  2HFI).  In  its  chemical  deportment 
it  is  an  acid  similar  to  the  hydrogen-halogen  acids.  Its  aqueous  solution 
reddens  blue  litmus-paper,  dissolves  many  metals,  and  saturates  bases,  form- 
ing salts  with  them,  in  which  two  hydrogen  atoms  are  replaced  by  metal. 

The  potassium  and  barium  salts  dissolve  with  difficulty  in  water. 


Fig.  63. 


Silicon  Carbide,  Carborundum,  SiC.  This  remarkable  substance  was  discovered  by 
Acheson  when,  at  the  suggestion  of  Edison,  he  endeavored  to  dissolve  carbon  in  fused  and 
highly  heated  aluminium  silicate,  hoping  that  when  the  mass  cooled  diamonds  would  appear. 
Miihlhauser  determined  its  composition.  Carborundum  is  produced  by  heating  a mixture 
of  sand  and  coke,  with  salt  as  a flux,  in  the  electric  furnace  to  about  3500°  : SiO.^  -f-  3C 
= SiC  -f-  2CO.  When  pure  it  consists  of  brilliant  green  crystals  of  specific  gravity  3.22. 
It  is  insoluble  in  the  ordinary  solvents  and  burns  with  difficulty  even  in  oxygen.  It  is 
attacked  by  molten  caustic  alkalies  and  alkaline  carbonates.  In  hardness  it  ranks  next 
to  boron  carbide  and  the  diamond  ; hence  it  is  beaten  into  powder  and  used  as  an  abra- 
sive, .superior  to  anything  known  at  present  in  this  line.  See  Metallic  Carbides  ; also  Z. 
f.  anorg.  Ch.  5 (1894I,  104;  O.  N.  Witt,  Die  chemische  Industrie  auf  der  Kolum- 
bischen  Weltaiusstellung  zu  Chikago  (1894),  124;  F.  Ahrens,  Die  Metallkarbide  und 
ihre  Verwendung,  Stuttgart,  1896. 


The  heats  of  for7ualio7i  of  the  halogen  derivatives  of  silicon  are  as  follows 

(Si,Cb):- 157.6;  (Si,I5r,)  = 120.4;  (Si,I,)  - 58.0. 


164 


INORGANIC  CHF.MISIRY. 


Their  ready  decoinposahility  l)y  water  may  he  aceomited  for  if  llie  lieat  f)f  formation 
of  silieic  acid  he  eonsidered  : Si,  O.^  ^ 219.0  Cal.  'I'lieir  depoilmeiil  toward  oxy{;en  is 

also  very  interesting.  The  heat  disengagement,  in  their  transposition,  following  the 
equation 

Six,  + 0,  = SiO.,  4-  4X, 

equals  61.4  Cal.,  with  the  ehloride,  98.6  Cal.,  with  the  hromide  and  with  the  iodide 
161.0  [=  219-58)  Cal.  Therefore,  oxygen  must  liberate  the  halogen  atoms  present  in 
them.  And  this  occurs  with  the  chloride  and  hromide  when  raised  to  a red  heat,  while 
the  iodide  inflames  in  the  air,  yielding  silicon  dioxide  and  iodine  vapor. 


Tin,  Sn  = 118.5,  recently  discovered  (1886)  Germanium, 

Ge  ==  72,  attach  themselves  in  chemical  character  to  the  groiij)  contain- 
ing carbon  and  silicon.  They  bear  the  same  relation  to  the  latter 
elements  which  arsenic  and  antimony  bear  to  the  members  of  the  nitrogen 
group.  This  is  evident  in  their  atomic  weights  : 

N = 14.04  P = 3i.o  As  =:  75  Sb  = 120 

C — 12.00  Si=28.4  Ge=72  80=118.5. 

Germanium  and  tin,  like  carbon  and  silicon,  form  volatile  compounds 
with  four  atoms  of  the  halogens,  e.  g.,  GeCl,  and  SnCl,.  In  a similar 
manner  they  unite  with  oxygen  to  yield  dioxides,  e.  g.,  GeOg  and  SnO^, 
which  resemble  silicon  dioxide  in  having  an  acid  character.  On  the 
other  hand,  they  do  not  form  volatile  hydrogen  derivatives,  and  are  thus 
markedly  distinguished  from  the  other  non-metals.  Therefore  germa- 
nium and  tin,  with  their  higher  analogue.  Lead,  will  be  treated  with  the 
metals. 


In  the  preceding  pages  we  have  considered  four  groups  of  elements, 
comprising  all  the  so-called  metalloids  (with  the  exception  of  boron). 
In  each  group  the  last  members,  possessing  the  highest  atomic  weights, 
exhibit  very  distinct  metallic  properties.  This  is  clearly  the  case  with  ger- 
manium and  tin,  antimony  and  arsenic  in  the  free  condition.  Tellurium 
and  selenium  also  (in  the  crystalline  modification)  possess  marked  metallic 
appearance;  finally,  iodine  has  a metallic  luster.  The  affinity  for  hydro- 
gen, on  the  other  hand,  diminishes  with  increase  in  metallic  character  ; 
the  hydrides  of  iodine,  tellurium,  antimony,  and  arsenic  are  very  un- 
stable, and  decompose  readily  into  their  constituents ; finally,  ger- 
manium, tin,  and  bismuth  do  not  combine  with  hydrogen. 

The  remarkable  relations  between  the  atomic  weights  of  the  elements 
of  the  four  groups  are  noticeable  in  the  following  table: 


C = 12.00 

N = 14.04 

0 = 16.00 

FI  = 19 

Si  = 28.4 

P = 31.00 

S = 32.06 

Cl  = 35-45 

(}e=  72 

As=  75 

Se  = 79.1 

Br  = 79.96 

Sn  = 1 18.5 

Sh  = 120 

Te  = 127 

I = 126.85. 

A fuller  consideration  of  these  relations  will  be  given  in  the  presenta- 
tion of  the  Periodic  System  of  the  Elements. 


THE  VALENCE  OF  THE  ELEMENTS. 


165 


THE  VALENCE  OF  THE  ELEMENTS.  CHEMICAL 
STRUCTURE  OF  THE  MOLECULES. 

Each  of  the  elements  of  the  group  of  halogens  forms  but  one  compound 
with  hydrogen  : the  halogen  hydride.  In  the  case  of  the  elements  of  the 
other  three  groups  just  considered  this  simplicity  occurs  exceptionally. 
Thus,  oxygen  and  sulphur  each  form  two  hydrogen  derivatives : 

H2O  and  H2O2 ; H^S  and  H^Sg. 

The  hydrides  are  even  more  numerous  with  phosphorus  and  nitrogen,  and 
with  carbon  they  are  almost  too  numerous  to  mention.  In  this  confusing 
multiplicity  simple  relations  can  be  reached  if  attention  is  directed  solely 
to  that  derivative  of  the  metalloid  which  contains  the  most  hydrogen  — 
to  the  compound  in  which  the  power  of  the  metalloid  to  unite  with  hydro- 
gen attains  its  maximum.  We  would  then  have  the  following  hydrides, 
which  in  turn  are  arranged  in  groups : 


CH4 

NH3 

OH2 

FIH 

SiH^ 

PH3 

SH2 

CIH 

ASH3 

SeH2 

BrH 

SbH3 

TeH2 

IH. 

Not  one  of  the  many  hundred  combinations  of  carbon  with  hydrogen 
contains  more  of  the  latter  element  than  methane;  ethane,  C2Hg,  follows 
methane  in  this  respect,  but  it  contains  only  three  atoms  of  hydrogen  to 
one  of  carbon,  and  in  all  the  other  hydrides  of  carbon  less  even  than  three 
atoms  of  hydrogen  are  present  for  each  atom  of  carbon.  It  must,  there- 
fore, be  concluded  that  the  carbon  atom  under  the  conditions  accessible 
and  known  to  us  is  not  capable  of  combining  with  more  than  four  atoms 
of  hydrogen.  The  same  is  true  of  the  other  elements  which  have  been 
mentioned.  Facts  oblige  us,  therefore,  to  ascribe  a peculiar  function  of 
affinity  to  each  element,  in  its  relation  to  hydrogen,  which  it  does  not 
exceed.  This  is  called  its  valence  or  ato7nicity.  The  elements  of  the 
fluorine  group  are  univalent;  the  elements  of  the  oxygen  group  bivakfit ; 
nitrogen  and  its  analogues,  U'ivalent ; carbon  and  silicon,  finally,  are 
quadrivalent  elements.  The  lary;est  number  of  hydro ge7i  atoins  with  which 
one  atoni  of  an  element  can  combine  determines  the  valence  of  the  latter.  It 
has  been  customary  to  attribute  to  the  elements,  in  accordance  with  their 
valency,  affinity  units  or  valence  units,  i.  e.,  to  say  that  carbon  has 
four  affinity  units,  oxygen  two,  etc.  This  is  intended,  however,  only 
to  indicate  that  one  carbon  atom  can  combine  with  but  four  hydro- 
gen atoms,  one  oxygen  atom  with  but  two  hydrogen  atoms,  etc.  (see 

p.  .66). 

The  valence  of  the  elements  is  frequently  designated  bylines  or  Roman 
numerals  placed  above  the  symbols  : 

I II  HI  IV 

Cl  O N C. 


i66 


INORGANIC  CHEMISTRY. 


The  mutual  union  of  affinity  units  is  indicated  by  one  line: 


II— Cl 

Hydroeen 

chloride. 


Water. 


II 

I 

N 

/ \ 

II  II 
Ammonia. 


II 

I 

II— C— II. 

I 

H 

Methane. 


In  these  formulas  the  atoms  of  oxygen,  of  nitrogen,  of  carbon — indeed, 
of  all  the  multivalent  elements — represent,  as  it  were,  the  nuclei  to  which 
the  hydrogen  atoms  attach  themselves. 

'I'he  hydrogen  atoms  in  these  molecules  can  be  replaced  or  substituted 
by  other  elements  (p.  159).  According  to  this,  an  atom  of  each  of  the 
univalent  halogens  can  replace  one  hydrogen  atom: 


II  u 
«<H 

II  II 

o<ci 

II  pi  III  /II 

o<ci 

III  /Cl 

Sb^Cl 

\ci. 

Water. 

Hypochlorous 

Chlorine  Nitrogen  Antimony 

acid. 

oxide.  iodide. 

trichloride. 

It  is  more  convenient  to  employ  brackets  instead  of  the  lines: 

1 

fll 

1 

fCl 

fll 

'c 

Cl 

'c 

Cl 

IV  Cl 

1 

1 

Cl 

1 

Cl 

1 

Cl 

[Cl. 

Chloroform. 

Carbon 

Silicon 

tetrachloride. 

chloroform. 

By  the  replacement  of  hydrogen  in  the  preceding  hydrogen  compounds 
by  the  univalent  metal  potassium,  we  obtain 


II  H 

0<K 

Ill  fH 

IV  1 

I H 

1 H 

K— Cl 

Potassium  chloride. 

K 

Ik 

Potassium  hydroxide. 

Dipotassium 

imide. 

[K. 

Potassium  methide. 

The  bivalent  elements,  like  oxygen  and  sulphur,  replace  two  hydrogen 
atoms  in  the  compounds  of  the  latter  element : 

III  /Cl 
Sb=0 

Antimony 
oxychloride 


Finally,  tri valent  nitrogen  can  replace  three  atoms  of  hydrogen,  or 
three  of  the  halogens  : 

III 

IV  f N I IV  III 

C\II  or  II-C=N. 

Ilydroj^en  cyanide. 

Ill 

IV  f N I IV  III 

C\C1  or  C1-C“N. 

Cyanogen  chloride. 


IV  /H 
C^H 

Methylene 

oxide. 


IV^O 

Carbon 

dioxide. 


iv^S 

Carbon 

bisulphide. 


IV^O 

s<o. 

Silicon 

dioxide. 


THE  VALENCE  OF  THE  ELEMENTS. 


167 


In  these  substitutions,  in  which  the  elements  replace  one  another  atom  for  atom,  we 
possess  a means  of  determining  the  molecular  weights,  previously  found  from  the  densities, 
in  a purely  chemical  manner.  Thus,  methane,  the  molecule  of  which  must  from  its 
density  be  represented  by  CH^,  can  have  its  hydrogen  atoms  replaced  one  after  the  other 
by  chlorine  and  four  compounds,  CH3CI,  CH2CI2,  CHCI3  and  CCl^  result.  Analysis  has 
determined  their  composition.  If  the  molecule  of  methane  were  represented  by  a larger 
formula,  say  CjHg,  then  upon  chlorination  there  would  result  not  four  but  eight  different 
chlorinated  methanes  : CgH^Cl,  CjHgCl.^,  C2H5CI3,  etc.,  of  which  four  (C2H.JCI,  C2HgCl3, 
C2H3CI5,  C2HCIY)  would,  so  far  as  composition  is  concerned,  differ  very  decidedly  from 
the  preceding.  Since,  however,  never  more  than  four  chlorinated  methanes  have  been 
made,  and  never  one  which  contained  more  hydrogen  than  CH3CI  or  one  having  less 
hydrogen  than  CIICI3,  we  conclude  in  a purely  chemical  way  that  methane  must  have  the 
formula  CH^.  This  method  of  determining  molecular  weights  can  also  be  applied  where 
the  vapor  density  is  not  applicable. 

These  substitutions  (replacements)  afford  a clue  to  the  magnitude  of  the  atomic 
weights  of  many  elements,  in  so  far  as,  as  has  already  been  repeatedly  emphasized,  the 
atouiic  tveight  is  the  smallest  quantity  of  an  element  which  is  present  in  the  moleczile  of 
any  of  its  compounds.  If  it  is  found  that  the  smallest  quantity  by  weight  of  chlorine 
which  enters  the  molecule  of  any  of  its  derivatives  is  to  the  replaced  hydrogen  as 
35.45  : 1. 01  (O  = 16)  then  it  must  be  assumed  that  these  numbers  represent  the  atomic 
weights  of  the  two  elements. 

Like  valence  is  designated  by  the  word  equivalence.  One  atom  of  chlorine  is  equiva- 
lent to  one  atom  of  hydrogen  ; 35.45  parts,  by  weight,  of  chlorine  are,  then,  equivalent 
to  1. 01  parts,  by  weight,  of  hydrogen.  One  atom  of  oxygen  is  equivalent  to  two  atoms 
of  hydrogen  ; consequently,  2.02  parts,  by  weight,  of  hydrogen  are  equivalent  to  16.00 
parts,  by  weight,  of  oxygen,  or  i.oi  parts  of  hydrogen  to  8 parts  of  oxygen.  Further, 
one  atom  of  nitrogen,  or  14.04  parts,  is  equivalent  to  three  atoms  or  3.03  parts  of 


hydrogen  ; i.oi  parts  of  hydrogen  is,  therefore,  equivalent  to  = 4.68  parts  of  nitro- 
gen, etc.  These  quantities,  equal  to  8 parts,  by  weight,  of  oxygen,  are  termed  equiva- 
lent weights y and  were  formerly  employed  instead  of  the  atomic  weights  (p.  70).  As 
may  be  observed  from  the  preceding,  the  equivalent  weights  of  multivalent  elements 
are  parts  of  the  atomic  weights  which  correspond  to  the  valence  of  the  elements. 


If,  consequently,  the  valence  of  the  elements,  in  relation  to  hydrogen 
(also  to  other  elements),  has  a definite  value,  the  question  naturally 
arises — what  will  result  if  an  atom  of  hydrogen  be  withdrawn  from  the 
saturated  molecules,  e.  g.,  water,  H2O,  ammonia,  NH3,  or  methane,  CH^? 
The  resulting  groups  or  residues : 

II  III  IV 

-O-H  -N^H2 

Hydroxyl.  Amide.  Methyl. 


can  plainly  not  exist  in  a free  condition.  When  set  free,  these  groups, 
therefore  (like  the  elementary  atoms),  unite  and  form  saturated  molecules. 
Thus,  for  example,  we  obtain  the  bodies: 


II  II  III  III 

HO-OH  H2N-NII2 

Hydrogen  Diamide, 

peroxide. 


Ill  III 


H2P-PH2 

Liquid 

hydrogen 

phosphide. 


IV  IV 
H3C-CH3. 
Dimethyl 
or 

Ethane. 


Carbon  is  particularly  inclined  to  such  combination.  Upon  remov- 
ing an  atom  of  hydrogen  from  dimethyl  or  ethane  (C2Hg)  the  so-called 
ethyl  group  remains : 

IV  IV 

C,U,-  or  CH3-CII2-. 


68 


INORGANIC  CHEMISTRY. 


in  wliicli  one  carlion  affinity  is  nnsatnrated  ; this  can  again  unite  with  llie 
methyl  gronj),  CII,.  'hhc  resnlting  com})onnd  would  be: 

IV  IV  IV 

cyig  or  HaC-CII^-CIIa  or  CjHg.CIIj. 

By  the  continuation  of  this  process  of  a chain-like  union,  as  it  were, 
of  the  carbon  atoms,  we  obtain  a whole  series  of  hydrocarbons  (CJ I, 
etc.)  with  the  general  formula  (compare  ])age  153). 

Not  only  similar  residues  or  groups,  but  dissimilar  also,  comlnne  in 
this  way : 

III  II  IV  II  IV  HI 

H^N-OII  IIjC-OII  II3C-NII2. 

Hydroxylamine.  Methyl  hydroxide.  Methylamiiie. 

Such  combinations  are  generally  effected  by  reactions  of  double  decomposition.  'Fluis 
methyl  hydroxide  (wood-spirit)  results  from  the  action  of  methyl  iodide  (CII3I)  upon 
silver  hydroxide  (AgOII)  : 

CII3I  -f  AgOII  = CII3 . Oil  b-  Agl  ; 

methylamine  by  the  action  of  methyl  iodide  upon  ammonia  : 

CH3I  -f  NII3  CII3  . NII2  + III. 

Dimethyl  is  produced  when  sodium  acts  upon  methyl  iodide : 

2CH3T  + Na^  = CJIg  + 2NaI. 

By  the  withdrawal  of  iodine  by  the  sodium  the  methyl  groups  are  liberated,  and  then 
combine  with  each  other. 

Further,  the  atoms  with  several  valences  can  unite  with  two  and  three 
affinities  {liouble  and  triple  union  ; there  being  no  intention  to  assert  with 
what  firmness  this  union  exists)  : 

III  III  IV  IV  IV  IV 

N^N  H2C=CH2  HC=CH. 

Ethylene.  Acetylene. 

o 

Hyponitrous 

oxide. 


The  molecules  of  elements  result  by  the  complete  mutual  union  of 
their  atoms : 


H-H  0-^0 

Hydrogen.  Oxygen. 


N=N 

Nitrogen. 


P-P 

P^P. 

Phosphorus. 


This  manner  of  linking  or  combination  of  the  atoms  in  the  molecule 
with  single  affinities  is  designated  cheinical  co?istitution,  or  chemical 
structure  oi  compounds;  the  formulas  representing  them  are  called  con- 
stitution or  structural  formulas.  Of  course,  the  actual  ])osition  of  the 
atoms  in  space  (of  which  we  have  no  knowledge)  is  not  indicated  by  the 
chemical  structure.  The  fundamental  ])rinci])le  of  cliemical  structure 
consists  in  this,  that  the  affinity  unit  of  one  atom  unites  with  the  affinity 
unit  of  another  atom. 

'I'he  following  circumstances,  however,  complicate  these  simple  rela- 


THE  VALENCE  OF  THE  ELEMENTS. 


69 


tions:  Among  the  elements  of  the  nitrogen  group  we  saw  that  phos- 
phorus and  antimony  combine  with  3 and  5 atoms  of  chlorine  and  the 
other  halogens;  that  sulphur,  selenium,  and  tellurium  take  up  2 and  4 
atoms  of  chlorine  and  bromine;  that  iodine  unites  with  i and  3 atoms 
of  chlorine  and  5 of  fluorine.  Only  the  quadrivalent  elements,  carbon 
and  silicon,  are  capable  of  combining  with  4 hydrogen  atoms,  and  with 
not  more  than  4 of  the  halogens : 


IV 

Ill 

II 

I 

CCl, 

PC13 

SC12 

ICl 

V 

IV 

III 

• • 

PCI, 

SCfi 

IC13 

V 

. . 

. . 

. . 

IF15. 

Hence  it  appears  that  the  elements  (excepting  carbon  and  silicon)  do  not 
express  such  a constant  valence  in  their  relation  to  chlorine  (and  to  the 
halogens)  as  they  do  to  hydrogen.  Phosphorus  and  its  analogues  appear 
to  be  tri-  and  quinquivalent ; the  elements  of  the  sulphur  group,  bi-  and 
quadrivalent;  iodine  finally  appears  to  be  uni-,  tri-,  and  quinquivalent. 

The  varying  valence  of  the  metalloids  shows  itself  more  distinctly  and 
frequently  in  the  more  stable  oxygen  derivatives.  We  are  acquainted 
with  the  following  oxygen  compounds  of  the  elements  of  the  four  groups 
already  mentioned ; the  members  of  each  group  afford  perfectly  analo- 
gous compounds  : 


II 

I 

III 

I 

CO 

N20 

S-^Og 

C120 

IV 

II 

IV 

IV 

C02 

NO 

S02 

C102 

III 

VI 

V 

N2O3 

SO3 

I205. 

IV 

VII 

. . 

N02 

S20, 

. . 

V 

. . 

N,05 

. . 

The  valence  of  iodine  and  of  nitrogen  reaches  five,  and  that  of  sulphur 
seven  affinity  units,  assuming  that  oxygen  is  constantly  bivalent  and 
wholly  united  to  the  other  element.  The  elements  of  the  nitrogen  group 
are  not  more  than  quinquivalent,  both  in  their  relation  to  the  halogens 
and  oxygen : 

PCI,,  PF1„  SbCl,. 

Carbon,  finally,  does  not  show  more  than  four  affinities  for  hydrogen, 
chlorine,  or  oxygen.  As  we  will  observe  later,  these  relations  are  more 
apparent  in  the  hydroxyl  derivatives  of  the  oxides — the  acids. 

Hence  we  conclude  that  valence  is  not  an  absolute  property  belonging 
per  se  to  and  for  the  elements  solely,  but  that  it  is  dependent  upon  the 
nature  of  the  elements  combining  with  one  another  and  upon  the  exter- 
nal conditions  under  which  they  unite.  We  can,  in  general,  distinguish 
two  valences:  the  hydrogen  valence  and  the  halogen  or  oxygen  valence. 
The  hydrogen  valence  is  constant  ior  all  elements  ; for  Cl  = i,  for  0 = 2, 
for  N = 3,  for  C = 4.  The  valence  of  most  of  the  elements  appears 
*5 


INORGANIC  CHEMISTRY. 


I 70 

to  be  variable  in  respect  to  oxygen  and  chlorine.  It  varies,  indeed,  as 
seen  from  the  formulas  given  above  for  the  chlorine  group  from  i to  4 to 
5 ; for  the  elements  of  the  sulphur  group  from  4 to  6,  and  from  3 to  7 ; 
for  those  of  the  nitrogen  group  from  i to  3 to  5,  and  from  2 to  4. 

It  is  perfectly  immaterial  whether  we  ascribe  a varialile  valence  to  the 
elements  or  accept  the  maximum  as  the  true  measure,  and  regard  the 
lower  compounds  as  unsafurated,  because  we  possess  no  concei)tion  of 
the  nature  of  valence.  Later,  we  will  discover  that  this  alteration  of 
valence  finds  full  ex})ression  and  generalization  in  the  periodic  system  of 
the  elements,  which  is  based  upon  the  grouping  of  the  elements  according 
to  their  atomic  weights. 


From  these  facts,  we  have  evolved  the  idea  of  a variable  valence.  Until  recently 
another  opinion  prevailed.  This,  denying  the  alterability,  regards  valence  as  an  abso- 
lute constant  property  of  the  elementary  atoms.  According  to  this  idea,  the  true  valence, 
or  atomicity,  is  only  derivable  from  the  hydrogen  comjjounds  ; the  halogens  are  absolutely 
univalent,  the  elements  of  the  oxygen  group  bivalent,  and  those  of  the  nitrogen  group 
trivalent,  etc.  Different  propositions  have  been  made  in  order  to  carry  out  the  constant 
atomicity  for  all  compounds.  A chain-like  union  is  assumed  in  order  to  explain  the 
OKygen  compounds.  This  is  similar  to  the  carbon  atoms  in  the  carbon  compounds 
(p.  168).  It  is  represented  in  the  following  examples: 


I II  II  II  II  II  I I II  II  II  II  I 

I_0— O— O— O— O— I Cl— O— O— O— ( )— H 

Iodic  anhydride.  Perchloric  acid. 


Sulphur  dioxide. 


II /O 

s(  )0 

Sulphur  trioxide. 


II  o— O— O— H 

^O— H 

Sulphuric  acid. 


II  III  II  II  II  III  II 

o N— O— O— O— N=0 

Nitrogen  pentoxide. 


II  III  II  II  I 
O N— O— O— H. 

Nitric  acid. 


Hence  it  would  appear  that  the  oxygen  atoms  can  unite  without  end,  in  a chain-like 
manner,  similar  to  the  carbon  atoms  in  the  carbon  compounds.  However,  a maximum 
combining  affinity  of  different  groups  for  oxygen  does  actually  occur  ; this  varies  regularly 
from  group  to  group  (for  iodine  = 5,  for  sulphur  ^ 7,  for  nitrogen  = 5,  for  carbon  =4). 
Consider  it  as  we  may,  on  the  supposition  of  constant  affinity,  the  reason  for  the  different 
tmmber  of  linking  oxygen  atoms  must  be  based  on  the  nature  of  the  other  element.  The 
fact,  then,  that  the  higher  oxides  or  their  hydrates  (HCIO4,  H2SO4,  HNO3)  are  more 
stable  than  the  lower  ( p.  182),  argues  decidedly  against  the  chain-like  union  of  the 
oxygen  atoms,  which  occurs  in  the  unstable  peroxides. 

To  explain  the  other  compounds  according  to  the  constant  atomicity  theory,  a differ- 
ence is  conceived  between  atomic  and  molecular  compounds.  The  former  are  such  as  can 
be  explained  by  constant  atomicity.  All  others  are  regarded  as  molecular  compounds, 
resulting  from  the  union  of  two  or  more  molecules,  upon  the  basis  of  newly  acquired 
molecular  affinities.  Thus  the  compounds  PCI5,  SCI4,  ICI3  are  viewed  as  addition  prod- 
ucts of  atomic  compounds  with  chlorine  molecules  : 

VC\,Cl,  SCl2,Cl2  ICfiCb- 

Phosphorus  Sulphur  Iodine 

peutachloridc.  lelrachloridc.  trichloride. 


THE  VALENCE  OF  THE  ELEMENTS. 


171 

It  was  thought  that  a proof  of  these  bodies  being  differently  constituted  from  the  real 
atomic  compounds  was  the  fact  that  when  vaporized  they  decomposed  into  simpler 
derivatives  ; the  molecular  compounds  were  not  regarded  as  capable  of  existing  as 
such  in  the  gaseous  state.  We  saw,  however,  that  the  decomposition  of  the  molecules 
PCI5  and  SCl^  is  only  gradual,  increasing  with  the  temperature,  and  that  they  do  exist  at 
lower  temperatures,  undecomposed,  in  a vapor  form  (compare  p.  141).  True  atomic 
compounds,  like  sulphuric  acid  and  nitric  acid,  frequently  separate  into  simpler  molecules, 
in  their  conversion  into  vapor  (compare  Sulphuric  Acid). 

Phosphorus  pentafluoride,  PFI5,  a gas  at  ordinary  temperatures,  is  a strong  argument 
in  favor  of  the  quinquivalence  of  phosphorus  ; iodine,  too,  forms  a volatile  pentafluoride, 
IFI5.  It  is  noteworthy  that,  as  a general  thing,  the  metalloids  yield  more  stable  and 
higher  compounds  with  the  lower  halogens  (fluorine  and  chlorine)  than  with  bromine 
and  iodine,  which  possess  higher  atomic  weights  (p,  142).  The  idea  that  molecular  com- 
pounds can  also  exist  in  vapor  form,  would  render  their  distinction  from  the  atomic  com- 
pounds purely  arbitrary — not  present  in  the  nature  of  the  bodies  themselves.  But  other 
gaseous  compounds  exist,  which  can,  in  no  light,  be  regarded  as  molecular.  Thus,  the 
usually  quadrivalent  or  sexivalent  tungsten  (WClg,  WBrg,  WOCI4)  forms  a gaseous  penta- 
chloride,  WCI5,  and  molybdenum,  perfectly  analogous  to  tungsten,  yields  a pentachloride, 
M0CI5.  Further,  the  quinquivalent  vanadium  (VdOClg)  yields  a gaseous  tetrachloride, 
VdCl^. 

The  salts  of  ammonia,  according  to  the  theory  of  constant  atomicity,  are  not  regarded 
as  ammonium  compounds : 

V V V 

NH,C1,  NH^NOg,  (NHJgSO, 

(p.  129),  but  as  addition  products  of  ammonia  with  acids  : 

NH3,HC1,  NH3,HNOg,  (NH3)3,H3SO„ 

which  would  make  the  analogy  of  the  same  with  metallic  salts  appear  rather  doubtful. 
Further,  the  properties  of  many  compounds  like 

V V IV  VI  V 

POCI3,  (C3Hg)3PO,  (CH3)3S.0H,  (CH3)3S0,  (CH3),N.0H, 

and  others,  cannot  be  interpreted  by  the  constant  valence  theory.  The  existence  of 
potassium  permanganate,  KMnO^,  is  not  compatible  with  a constant  bi-  or  quadrivalence 
of  manganese  (see  p.  178). 

Yet,  down  to  the  present  time,  the  acceptance  of  molecular  additions  could  not  be  en- 
tirely dispensed  with,  especially  for  the  so  called  water  of  crystallization  compounds;  the 
proposed  effort,  however,  continues  to  reduce  all  such  compounds  on  the  basis  of  higher 
valence  of  the  elements.  While,  consequently,  the  constant  valence  theory  comprises 
only  the  so-called  atomic  compounds,  the  extended  valence  idea  draws  all  others  into  the 
circle  of  generalization. 

We  must,  first  of  all,  remember  that  the  nature  of  chemical  union  and  the  cause  of  the 
valence  of  the  atoms  are  entirely  unknown  to  us,  and  that,  therefore,  neither  the  idea  of 
a variable,  nor  yet  of  a constant  valence  constitutes  a final  explanation.  Our  task  must 
be  restricted  rather  to  representing  the  varying  combination  relations  of  atoms  in  accord- 
ance with  the  facts  and  free  them  as  far  as  possible  from  hypothetical  additions.  The 
supposition  of  a constant  valence,  formerly  given  preference  as  the  simpler,  has  shown 
itself  to  be  insufficient.  Because  it  maintains  real  differences,  where  such  are  not  percep- 
tible, it  departs  from  the  ground  of  induction.  The  idea  of  a variable  valence,  which 
takes  all  facts  into  consideration,  is  not  prejudicial.  It  is  their  simplest  expression. 
Indeed,  it  finds  a pregnant  analogy  in  the  deportment  of  the  hydrocarbon  radicals.  By 
the  elimination  of  hydrogen  from  the  saturated  molecules  (C2H6,  for  example),  we  obtain 
radicals  or  groups  of  increasing  valence  (CjHg,  C2.H3,  C2H2'),  The  group 

C2H2,  however,  is  bi-  and  quadrivalent;  the  group  Cgllg,  uni-  and  trivalent.  It  is  very 
probable  that  the  elementary  atoms  are,  like  these  radicals,  of  a compound  nature. 


172 


INORGANIC  CHEMISTRY. 


The  principles  of  chemical  structure,  just  presented,  appear  most 
clearly,  and  with  the  greatest  regularity,  with  carbon.  The  consti- 
tution of  the  innumerable  varieties  of  carbon  compounds  is  exjilained 
by  the  quadrivalent  nature  of  the  carbon  atoms,  and  their  ability  to  com- 
bine with  each  other  by  single  affinities.  In  the  other,  the  so-called 
inorganic  compounds,  the  valence  and  structure  relations  are  more 
complicated,  and  are  far  less  investigated,  but  even  in  them  so  many 
regularities  appear  that  the  actual  material  is  greatly  simplified  therebv, 
and  is  made  more  comprehensible.  The  doctrine  of  valence  and  struc- 
ture is  the  first  attempt  to  refer  the  facts  underlying  the  law  of  multiple 
proportions  to  the  properties  of  the  elementary  atoms.  As  this  theory 
merely  comprises  actual  relations,  it  cannot  be  negatived,  but  only  further 
developed. 


The  idea  of  varying  valence  of  the  atoms  lay  undeveloped  as  the  basis  of  Gerhardt’s 
type  theory,  according  to  which  chemical  compounds  were  referred  to  a few  fundamental 
forms  or  types  : 


H 

H 


11 

Cl 


H 

IIN. 

H 


This  mode  of  comparison  will  be  repeatedly  adopted  in  the  following  pages  (p.  198). 
After  Edward  Frankland  (1852)  had  established  the  doctrine  of  the  saturating  capacity 
of  the  elementary  atoms,  August  Kekule  added  the  type  CIb  to  those  just  given  and 
solved  in  a clear  and  convincing  manner  the  question  how  carbon  atoms  were  able  to 
combine  with  one  another  and  to  satisfy  their  affinities  (1857-1859).  In  this  way  he 
developed  one  of  the  fundamentals  of  the  chemistry  of  to-day— doctrine  of  the  linking 
of  the  atoms.  See  the  historical  sketch  in  Richter’s  Organic  Chemistry. 


OXYGEN  COMPOUNDS  OF  THE  METALLOIDS. 

Almost  all  the  oxygen  derivatives  of  the  metalloids  yield  acids  with 


water : 

H2O  = 2HIO3; 

Iodic 

SO3  -f  H2O 

= H2SO, ; 

Iodine 

Sulphur 

Sulphuric 

pentoxide. 

acid. 

trioxide. 

acid. 

P2O5  + 3^2^^ 

Phosphorus 

pentoxide. 

= 2H3PO4. 

Phosphoric 
acid. 

Conversely,  these  oxides  may  be  obtained  by  the  removal  of  water  from 
the  acids;  therefore,  they  are  ordinarily  termed  anhydrides  of  the  corre- 
sponding acids:  laO^,  anhydride  of  iodic  acid;  SO3,  sulphuric  anhydride; 
PjOj,  phosphoric  anhydride,  etc. 

Salts  are  formed  by  the  replacement  of  the  hydrogen  of  the  acids  by 
metals.  The  acids  are  distinguished  as  monobasic,  dibasic,  tribasic, 
and  polybasic,  de])ending  upon  the  number  of  their  hydrogen  atoms 
reidaceable  by  metals. 

Like  the  metalloids,  the  metals  possess  different  valences  ; the  univalent 
metals  (sodium,  potassium,  silver)  each  replace  one  hydrogen  atom; 


OXYGEN  COMPOUNDS  OF  THE  HALOGENS. 


173 


those  of  higher  valence  replace  more.  Therefore  the  metals  of  higher 
valence  can  unite  several  acid  residues.  This  explains  the  import,  e.  g., 
of  the  following  chemical  formulas  : 


(1)2 

K2SO, 

Potassium  sulphate. 
I 

KNO3 

Potassium  nitrate. 


I  I II 

KNaSO^  CaSO^ 

Potassium-sodium  sulphate.  Calcium  sulphate. 

II  III 

Cu(N03)2  Bi(N03)3. 

Copper  nitrate.  Bismuth  nitrate. 


The  salts  of  sulphuric  acid  are  called  sulphates,  those  of  nitric  acid, 
nitrates,  those  of  phosphoric  acid,  phosphates,  etc.  The  symbols  of  the 
metals  are  sometimes  written  at  the  beginning  and  sometimes  at  the  end 
of  the  chemical  formulas  of  the  salts ; in  the  second  case  we  mean  to  indi- 
cate that  the  metal  atoms  are  in  union  with  the  oxygen. 


1.  OXYGEN  COMPOUNDS  OF  THE  HALOGENS. 


Oxygen  derivatives  of  fluorine  are  not  known.  Iodine  is  the  only  one 
of  the  remaining  halogens  which  unites  directly  with  oxygen  ; how'- 
ever,  the  oxides  of  chlorine  may  be  prepared  by  indirect  methods.  The 
only  known  acids  of  bromine  are  those  corresponding  to  the  oxides. 

Chlorine  forms  the  following  anhydrides  and  acids  : 

Anhydrides.  Acids. 

CljO  HCIO,  Hypochlorous  acid. 

. . (HCIO2,  Chlorous  acid.) 

CIO2  or  CI2O, 

. . HCIO3,  Chloric  acid. 

. . HCIO4,  Perchloric  acid. 


The  compound  Cl20^  exists,  and  must  be  regarded  as  a mixed  anhy- 
dride of  chlorous  and  chloric  acids.  Chlorous  acid  itself  is  not  known 
free,  but  is  stable  in  the  form  of  salts.  Further,  the  anhydrides  of  chlorous, 
chloric  and  perchloric  acids,  having  the  formulas  CI2O3,  CljO^and  Cl20^, 
are  not  known. 

The  following  formulas  express  the  chemical  structure  of  these  com- 
pounds : 


I II  I 
Cl-O-Cl 

Hypochlorous 

anhydride. 


I II  I 
Cl-O-H. 
Hypochlorous 
acid. 


Ill  V 

OCI-O-CIO2 

Chlorous-chloric 

anhydride. 


V 


CIO2-OH. 

Chloric 

acid. 


VII 

CIO3-OH. 
Perchleric  acid. 


174 


INORGANIC  CHEMISTRY. 


In  tlic  acids  vve  assume  tlie  presence  of  the  univalent  group  OH 
(hydroxyl  or  water  residue)  the  hydrogen  of  which  can  be  replaced  by 
metals,  through  the  action  of  metals  or  bases.  The  grouj)  (CIO,^  or  CIO.,) 
combined  with  hydroxyl  is  called  an  acid  7'esidue,  or  radical.  In  the 
anhydrides  of  the  monobasic  acids  two  acid  radicals  are  united  by  an 
oxygen  atom;  water  converts  them  into  2 molecules  of  acid  : 


g>0  + 0<J| 


Cl-OII 

Cl-OII. 


The  salts  of  perchloric  acid  are  called  perchlorates ; those  of  chloric 
acid,  chlorates ; those  of  chlorous  acid,  chlorites ; and  those  of  hypo- 
chlorous  acid,  hypochlorites. 

Hypochlorous  Oxide,  Cl^O  (chlorine  monoxide,  hypochlorous  an- 
hydride), is  i)roduced  by  conducting  dry  chlorine  gas  over  ]jrecipitated 
mercuric  oxide,  cooled  by  ice  or  cold  water;  the  brown  oxychloride  of 
mercury  is  formed  at  the  same  time  : 


2HgO  + 2CI,  = IlgpCl,  + Cip. 

The  precipitated  mercuric  oxide  should  be  heated  to  300°  before  it  is  used,  otherwise 
it  will  be  acted  upon  too  energetically  by  the  chlorine,  and  oxygen  instead  of  hypochlo- 
rous oxide  will  result.  Red  mercuric  oxide  is  affected  too  slowly  by  chlorine. 

The  disengaged  gas  is  condensed  in  a bent  glass  tube,  cooled  by  a 
freezing  mixture  of  ice  and  calcium  chloride. 

It  is  a reddish-brown  liquid,  boiling  at  -j-5°.  In  the  gaseous  condition  it 
has  a yellow-brown  color,  an  extremely  disagreeable  odor,  and  it  attacks  the 
respiratory  organs  very  strongly.  The  vapor  density  corresponds  to  the 
molecular  formula,  CI2O  = 86  9.  The  gaseous  or  liquid  oxide  is  very 
unstable  and  explosive.  If  heated,  or  if  brought  in  contact  with  a flame, 
or  exposed  to  the  action  of  the  electric  spark,  it  decomposes  with  deto- 
nation into  oxygen  and  chlorine  : 

2CI.2O  = 02  + 2CI2. 

2 vols.  I vol.  2 vols. 


It  is  similarly  exploded  when  it  is  brought  in  contact  with  sulphur, 
phosphorus,  and  organic  substances,  such  as  caoutchouc.  It  has  strong 
oxidizing  and  bleaching  properties.  It  dissolves  in  water  to  hypo- 
chlorous acid : 

CI2O  + H2O  = 2IICIO. 


It  yields  calcium  hypochlorite  and  chlorine  when  it  is  conducted  over 
porous  calcium  chloride : 


CaCl2  -}-  2CI2O  = Ca(C10)2  + 2CI2. 


Hypochlorous  Acid,  HCK),  is  only  known  in  aqueous  solution. 
Its  presence  in  chlorine  water  has  been  mentioned  (p.  52).  It  is 
ol)tained  liy  conducting  clilorine  into  water  in  which  there  is  suspended 


HYPOCHLOROUS  ACID.  I 75 

freshly  precipitated  mercuric  oxide;  the  liquid  is  agitated  constantly, 
while  at  the  same  time  it  is  cooled  and  shut  olf  from  the  light : 

HgO  -1-  2CI2  -h  H2O  = HgCh  + 2HCIO. 

The  mercuric  chloride  produced  combines  with  the  excess  of  oxide  to 
form  oxychloride  (see  p.  174).  After  filtration  the  solution  may  be  dis- 
tilled ; a strong  chlorine  odor  is  usually  noticeable  at  the  beginning.  The 
first  distillates  are  especially  rich  in  hypochlorous  acid.  The  free  chlorine 
is  expelled  by  blowing  air  through  the  solution  (Balard). 

Hypochlorous  acid  may  be  similarly  prepared  by  conducting  chlorine 
into  water  in  which  chalk  is  suspended  : 

CaCOg  -j-  2CI2  T HjO  = CaCq  -f-  CO2  T 2HCIO  (Williamson)  ; 

also  by  the  distillation  of  bleaching  lime  solutions  with  as  much  dilute 
acid  as  is  required  for  the  decomposition  of  the  hypochlorite,  e.  g. : 

NaClO  -f  NaCl  + HNO3  = NaNOg  -f  NaCl  -f  HCIO  (Gay-Lussac). 

Boric  acid  is  well  adapted  for  this  purpose  as  it  does  not  decompose 
the  chloride.  The  concentrated  aqueous  solutions  (5  per  cent.)  are 
yellow  in  color.  Upon  distillation  there  ensues  a partial  decomposition 
into  chloric  acid  and  chlorine: 

{a)  3HCIO  = 2HCI  + HClOg ; 

\b)  2HCI  + 2HCIO  = 2CI2  + 2H2O. 

The  decomposition  is  very  slight  with  dilute  solutions.  These  are, 
unlike  the  concentrated  liquids,  quite  stable  in  the  light.  Their  odor  is 
peculiar  and  they  oxidize  and  bleach  energetically.  The  bleaching  action, 
due  to  the  separation  of  oxygen  in  statu  nascendi,  is  twice  as  great  as  that 
of  free  chlorine,  as  is  evident  from  the  following  equations: 


CI2  + H2O  = 2HCI  -f  O, 

2HCIO  = 2HCI  + O2. 

Hydrochloric  acid  decomposes  hypochlorous  acid  into  chlorine  and 
water,  when  just  twice  as  much  chlorine  is  evolved  as  there  is  present  in 
the  hypochlorous  acid. 

The  acid  itself  is  very  feeble,  more  feeble  than  carbonic  acid,  hence  it 
is  produced  free  if  chlorine  be  conducted  into  aqueous  solutions  of  car- 
bonates : 

K2COg  -f  CI2  -h  H2O  = KCl  -f-  KHCOg  -f  HCIO. 

Its  salts  (bleaching  powder,  see  Chloride  of  Lime)  are  formed  together 
with  chlorides  by  the  action  of  chlorine,  in  the  cold,  upon  strong  bases: 

2NaOH  + CI2  NaCl  -f  NaClO  + H2O. 

Within  certain  limits  this  transposition  takes  place  independently  of 
the  temperature  and  strength  of  solution  if  chlorine  is  not  present  in 
excess,  otherwise  the  formation  of  chlorates  will  occur  even  at  low  tem- 
peratures. When  the  solutions  are  heated  two  decompositions  of  the 


INORGANIC  CHEMISTRY. 


I 76 

hy})ochlorite  ])roceed  simultaneously;  in  one  instance  chloride  and 
chlorate  result : 

3NaC10  ^ 2NaCl  + NaClOj, 
and  in  the  other  case  chloride  and  oxygen  : 

2NaC10  = 2NaCl  + O,. 

[Compare  Foerster  and  Jorre,  Jr.  prakt.  Ch.  59  (1899),  53  I.] 

A hypochlorite  is  also  produced  when  chlorine  is  conducted  over  cal- 
cium hydrate  powder  (see  Bleaching  Lime). 

Oh  shaking  the  aqueous  solution  of  hypochlorous  acid  with  mercury,  there  is  produced 
a yellow-brown  precipitate  of  ligO.  llgCh,  soluble  in  hydrochloric  acid  (salts  of  hypo- 
chlorous  acid  form  HgO).  This  behavior  serves  to  distinguish  hypochlorous  acid  from 
chlorine,  which  under  like  circumstances  forms  mercurous  chloride,  Ilg^Ck,  insoluble  in 
hydrochloric  acid  (reaction  of  Wolters). 

Chlorine  Trioxide,  CI2O3  = (C10)20,  the  anhydride  of  chlorous 
acid  (p.  131),  and  the  free  Chlorous  Acid,  HClOj  = CIO. OH,  are 
unknown.  It  was  formerly  supposed  that  the  tri oxide  resulted  by  the 
reduction  of  chloric  acid,  if  potassium  chlorate,  KCIO3,  was  decomposed 
by  nitric  acid  in  the  presence  of  reducing  substances  (arsenic  trioxide, 
sugar  and  tartaric  acid).  Later  research  has,  however,  shown  that  in 
such  decompositions  it  is  not  the  trioxide,  but  the  dioxide,  CIO2,  mixed 
with  chlorine  and  oxygen,  which  results. 

The  salts  of  chlorous  acid,  the  chlorites,  e.  g. , potassium  chlorite,  are  formed  by  the 
transposition  of  the  aqueous  solution  of  chlorine  dioxide  with  the  alkalies  (see  below). 
The  silver  salt,  KgQXO^,  and  the  lead  salt,  Pb(C102).^,  are  precipitated  from  the  aqueous 
solutions  of  the  alkali  salts  by  silver  nitrate  and  lead  acetate.  They  dissolve  with  diffi- 
culty in  cold  water,  and  crystallize  from  hot  water  in  yellowish-green  scales.  They  are 
decomposed  by  carbonic  acid  ; the  liberated  chlorous  acid  decomposes  immediately. 

Chlorine  Dioxide,  CIO2,  or  Chlorine  Tetroxide,  CI2O4,  formerly 
called  hypochloric  acid,  is  formed  when  concentrated  sulphuric  acid  acts 
in  the  cold  upon  potassium  chlorate.  The  chloric  acid,  which  first 
appears,  at  once  decomposes  into  chlorine  dioxide,  water,  and  the  stable 
perchloric  acid  : 

3HCIO3  =z  CI2O,  -h  HCIO,  + HjO. 

It  is  best  obtained  (mixed  with  carbon  dioxide)  by  adding  dilute  sul- 
jdmric  acid  (with  2 parts  of  water)  to  a mixture  of  potassium  chlorate  (i 
])art)  and  oxalic  acid  (4.5  parts).  Upon  the  application  of  a gentle  heat 
a yellow-green  gas  escapes.  It  can  be  condensed  by  a freezing  mixture 
to  a reddish-brown  liquid,  boiling  at  9.9°  (under  a pressure  of  740  mm.), 
and  solidifying  at  — 79°.  This  liquid  may  be  distilled  without  decom- 
position from  glass  vessels,  free  from  organic  matter,  if  the  temperature 
does  not  exceed  30°.  Violent  explosions  occur  at  higher  temperatures, 
so  that  it  is  advisable,  if  ])roper  precautions  are  not  taken,  not  to  pour 
sulphuric  acid  U])on  ])otassium  chlorate.  Liquid  as  well  as  gaseous  chlo- 
rine-dioxide is  very  explosive,  especially  in  contact  with  organic  matter 
and  when  heated,  but  sunlight  is  without  effect. 

The  formation  of  the  oxide  can  be  effected  in  a perfectly  harmless  way. 


CHLORIC  ACID. 


177 


and  its  powerful  oxidizing  action  be  illustrated,  by  throwing  some  potas- 
sium chlorate  and  a few  pieces  of  yellow  phosphorus  into  water,  con- 
tained in  a measuring  glass,  then  allowing  sulphuric  acid  to  touch  the 
bottom  of  the  tube,  drop  by  drop,  by  means  of  a pipette.  By  the  action 
of  the  disengaged  dioxide  the  phosphorus  will  burn  under  water  with  a 
brilliant  light. 

When  concentrated  sulphuric  acid  is  added  to  a mixture  of  potassium 
chlorate  and  sugar,  a violent  combustion  occurs. 

Chlorine  dioxide  dissolves  in  water  with  a yellow  color.  Alkalies  de- 
colorize the  solution,  forming  salts  of  chloric  and  chlorous  acids ; 

no  + 2KOH  = CIO.OK  + CIO2.OK  -f  H^O. 

Potassium  Potassium 
chlorite.  chlorate. 

Sunlight  decomposes  the  aqueous  solution  of  chlorine  dioxide  into 
chloric  acid,  oxygen,  and  chlorine  : 

3CI2O4  + 2H2O  = 4CIO2.OH  + O2  4-  CI2. 

The  gas  density  of  the  dioxide  corresponds  to  the  molecular  formula 
CIO2.  It  is,  however,  very  probable  that  at  lower  temperatures,  in  the 
liquid  condition  or  in  aqueous  solution,  the  molecules  possess  the  doubled 
formula,  Cl20^,  and  that  then  the  compound  represents  the  mixed  anhy- 
dride of  chloric  and  chlorous  acids  : 

Cl  O — '>0 

- ci02->^- 

The  decomposition  of  its  aqueous  solution  by  alkalies  is  an  argument 
in  favor  of  this  view,  as  is  also  its  analogy  with  nitrogen  dioxide,  NOj, 
or  nitrogen  tetroxide,  N2O4  (see  this),  the  existence  of  both  of  which 
molecules  has  been  proved. 

Chloric  Acid,  HCIO3  or  CIO2.OH,  is  obtained  by  decomposing 
an  aqueous  solution  of  barium  chlorate  with  sulphuric  acid : 

Ba(C103)2  + H2SO,  = BaSO,  + 2HCIO3. 

Barium  chlorate.  Barium 

sulphate. 

The  barium  sulphate  separates  as  a white,  insoluble  powder,  and  can 
then  be  filtered  off  from  the  aqueous  solution  of  the  acid.  This  is  con- 
centrated without  decomposition,  under  diminished  pressure,  until  the  spe- 
cific gravity  becomes  1.28,  and  it  then  contains  about  40  per  cent,  of 
chloric  acid ; it  is  oily  and,  when  heated  to  40°,  decomposes  into  chlorine, 
oxygen,  and  perchloric  acid,  HCIO^.  The  concentrated  aqueous  solu- 
tion oxidizes  strongly;  sulphur,  phosphorus,  alcohol,  and  paper  are 
inflamed  by  it.  Hydrochloric  acid  eliminates  chlorine  from  the  acid  and 
its  salts: 

HCIO3  + SHCI  = 3II2O  + 3CI2. 

The  chlorates  are  slowly  formed  when  aqueous  hypochlorite  solutions 
are  heated  ; the  latter  partly  break  down  into  oxygen  and  chlorides.  To 


178 


INORGANIC  CHEMISTRY. 


obtain  the  chlorate  the  liypochlorite  solution  should  be  slightly  super- 
saturated with  chlorine;  then  the  chlorate  will  form  even  in  the  cold  but 
much  more  rapidly  if  the  solution  be  heated  to  130°.  This  may  be  thus 
explained.  The  excessive  chlorine  liberates  hy^jochlorous  acid  from  the 
hypochlorite : 

2KCIO  + 2CI2  f 211,0  = 2KCI  + 4IICIO, 


which  in  turn  is  rapidly  and  completely  trans])osed  by  the  hypochlorite 
into  chlorate  when  chlorine  or  hydrochloric  acid  is  again  liberated : 


or 


2KCIO  T 2IICIO  = KCIO3  + KCl  + 11,0  + Cl,, 
KCIO  + 2IICIO  = KCIO3  + 2IICI ; 


and  hypochlorous  acid  is  evolved  by  the  liberated  chlorine  or  hydro- 
chloric acid,  etc.  The  statement  previously  made  that  chlorine  acting 
upon  cold  lyes  produces  hypochlorites,  and  chlorates  when  the  solutions 
are  hot,  does  not  harmonize  with  the  facts.  The  chlorate  formation  is 
rapid  and  complete  only  when  chlorine  is  in  excess  (compare  Foerster 
and  Jorre).  The  electrolytic  production  of  chlorates  from  chlorides  will 
be  described  in  connection  with  potassium  chlorate. 

Perchloric  Acid,  HCIO^  or  CIO3 . OH.  This  is  the  most  stable 
of  all  the  oxygen  derivatives  of  chlorine.  Its  sodium  salt  is  present  in 
Chile  saltpeter.  As  previously  stated,  it  is  produced  by  the  decomposi- 
tion of  chloric  acid,  but  is  more  easily  obtained  from  its  salts.  Upon 
heating  potassium  chlorate  to  fusion,  oxygen  escapes  and  potassium  per- 
chlorate results  : 

2KCIO3  = KCIO4  -f  KCl  -f  O,. 

When  this  decomposition  approaches  completion,  the  fused  mass  becomes 
a thick  liquid  and  finally  a solid.  As  potassium  perchlorate  dissolves  with 
difficulty  in  water  it  can  be  readily  separated  from  the  chloride  formed 
simultaneously. 

A solution  of  perchloric  acid  containing  a little  sodium  chloride,  but  applicable  for  all 
analytical  purposes,  may  be  prepared  as  follows  : Commercial  sodium  chlorate  is  con- 
verted by  heat  into  perchlorate  and  chloride.  Concentrated  hydrochloric  acid  is  poured 
upon  the  powdered  mixture  and  dissolves  almost  nothing  but  perchloric  acid  : 

NaClO,  + HCl  = NaCl  + HCIO^. 

The  solution  is  filtered  and  evaporated  until  dense  white  fumes  of  perchloric  acid  com- 
mence to  escape  ; by  distillation  under  greatly  reduced  pressure  the  dihydrate  may  be 
obtained  [Z.  f.  anorg.  Ch.  9 (1895),  342;  13  (1897),  166]. 

Roscoe  obtained  a monohydrate  of  perchloric  acid  by  distilling  potas- 
sium perchlorate  with  four  times  its  quantity  of  concentrated  sulphuric 
acid  until  the  drops,  passing  over,  no  longer  solidified.  The  mass  was 
then  heated  to  110°  until  crystals  a])peared  in  the  neck  of  the  small 
retort.  From  this  point  on  anhydrous  perchloric  acid  distilled  over. 

The  pure  acid  is  a mobile,  colorless  liquid,  fuming  strongly  in  the  air; 
its  specific  gravity  is  1.78  at  15°.  It  cannot  be  preserved  even  in  the 
dark,  since  after  a few  days  it  decomposes  with  violent  explosion.  Heat 
jiroduces  the  same  result.  It  also  explodes  in  contact  with  phosphorus, 


BROMIC  ACID. 


179 


paper,  carbon,  and  other  organic  substances.  It  produces  painful 
wounds  when  brought  in  contact  with  the  skin.  It  absorbs  water  with 
avidity,  and  with  one  molecule  of  the  same  forms  the  crystalline  hydrate 
HClO^-f  H2O,  melting  at  50°  and  passing  at  110°  into  the  anhydrous 
acid  and  the  dihydrate,  HCIO^  + 2H2O  : 

2C10,H  . H2O  = HCIO,  + HCIO^  2H2O. 

The  dihydrate  may  also  be  obtained  by  evaporating  the  aqueous  solutions 
of  perchloric  and  chloric  acids.  It  is  a colorless,  oily  liquid  of  com- 
paratively great  stability.  It  boils  unchanged  at  203°.  Its  specific  gravity 
equals  1.82. 

Perchloric  acid  is  also  formed  at  the  positive  pole  by  the  electrolysis 
of  aqueous  solutions  of  inorganic  chlorine  compounds;  this  is  due  to  the 
liberation  of  oxygen  at  the  kathode. 


Bromine  yields  the  following  compounds  with  oxygen  and  hydrogen  : 

HBrO,  Hypobromous  acid. 

HBrOg,  Bromic  acid. 

HBrO^,  Perbromic  acid. 

The  corresponding  anhydrides  are  not  known.  The  acids  are  perfectly 
analogous  to  the  corresponding  chlorine  compounds. 

Hypobromous  Acid,  HBrO,  is  formed  when  bromine  water  acts 
upon  mercuric  oxide  ; the  aqueous  solution  can  be  distilled  in  vacuo,  and 
possesses  properties  very  similar  to  those  of  hypochlorous  acid. 

Bromic  Acid,  HBrOg.  Bromates  are  like  the  chlorates.  An 
aqueous  solution  of  the  acid  can  be  obtained  from  the  barium  salt  by 
decomposing  the  latter  with  sulphuric  acid.  A more  practical  method 
of  getting  the  free  acid  is  to  let  bromine  act  upon  silver  bromate : 

SAgBrOg  -f  6Br  -f  3H2O  = SAgBr  -f  fiHBrOs, 
or  to  oxidize  bromine  with  hypochlorous  acid  : 

5CI2O  + Br2  -f  H2O  2HBr03  + loCl. 

The  aqueous  solution  may  be  concentrated  in  vacuo  until  its  content 
reaches  50.6  per  cent.  HBrOg,  and  then  closely  corresponds  to  the 
formula  HBrO,  -f-  yH^O.  When  heated  it  decomposes  into  bromine, 
oxygen  and  water.  Heat  breaks  down  the  alkali  bromates  into  bromide 
and  oxygen  without  the  production  of  a perbromate. 

Perbromic  Acid,  HBrO^,  is  said  to  be  formed  in  the  action  of 
bromine  vapor  upon  perchloric  acid  : 

IICIO,  + Br  = HBrO,  -f  Cl, 
and  is  perfectly  similar  to  the  latter. 


i8o 


INORGANIC  CHEMISTRY. 


Iodine  forms  llie  following  anhydrides  and  acids; 

I-A  IIIO3  11104.211/) 

loclitie  pentoxide,  Iodic  acid.  Periodic  acid  (dihydrate). 

Iodic  anhydride. 


Iodic  Acid,  HIO3.  The  potassium  and  sodium  salts  of  this  acid 
occur  in  Chile  saltpeter  and  are  at  present  the  best  sources  for  iodine  (see 
p.  55).  The  iodates  are  formed  in  the  same  manner  as  the  chlorates  and 
bromates,  by  dissolving  iodine  in  a hot  solution  of  j)otassium  or  sodium 
hydroxide : 

6KOII-|-3l,  = 5KI  4-  KI()3  + 311.0. 


Upon  adding  barium  chloride  to  this  solution  sparingly  soluble  barium 
iodate  separates;  it  may  be  transposed  by  sulphuric  acid  into  insoluble 
barium  sul})hate  and  free  iodic  acid. 

When  iodine  acts  upon  cold  .sodium  or  potassium  hydrate  hypoiodites  are  formed. 
Schonbein  suspected  this,  but  Lonnes  proved  it  : 

2NaOII  + 2l  = NalO  + Nal  + Iip. 

These  salts,  however,  quickly  pass  into  iodates  : 

3XaIO  = NalOa  + ^Nal. 

Therefore  iodine  acts  upon  hydroxides  just  like  chlorine  and  bromine,  but  with  iodine 
the  second  step  of  the  reaction,  the  formation  of  iodate,  proceeds  very  rapidly  at  the 
ordinary  temperature.  The  experiments  of  Binz  and  also  those  of  Lonnes,  however, 
show  that  iodine  can  remain  uncombined  for  quite  a time  when  dissolved  in  alkaline 
liquors  if  potassium  iodide  be  present. 

Willgerodt  and  V.  Meyer  have  prepared  organic  derivatives  of  the  yet  unknown 
III 

H — I O.  Indeed,  Meyer  has  made  organic,  strongly  basic  compounds  which  may  be 

III 

derived  from  iodoniiim  hydroxide,  H2  = I — OH,  also  unknown.  The  chemical  struc- 
ture of  these  derivatives  forbids  the  assumption  that  iodine  is  invariably  univalent  (see 
Ber.  27  (1894),  426,  1592,  and  also  Richter’s  Organic  Chemistry). 

The  free  acid  can  be  obtained  by  the  oxidation  of  iodine  with  strong 
nitric  acid,  or  by  means  of  chlorine  : 

3I2  + loHNOj  = 6HIO3  + loNO  + 2H2O. 

Iodates  are  also  produced  by  the  action  of  iodine  upon  chlorates  or  bro- 
mates in  aqueous  solution,  whereby  the  iodine  directly  eliminates  the 
chlorine  and  bromine: 

KCIO3  + I = KIO3  + Cl. 

Upon  evaporating  the  aqueous  solution  the  free  iodic  acid  crystallizes  in 
colorless  rhombic  jirisins  of  sj)ecific  gravity  4.63.  When  iodic  acid  is 
heated  to  170°  it  decomposes  into  water  and  iodic  anhydride  : 

211103  = 1203  + 1130. 

It  is  decomposed,  like  chloric  acid,  by  hydrochloric  acid  : 

2IIIO3  + loIICl  = I2  + 5CI2  + 6II2O. 


PERIODIC  ACID. 


l8l 


Reagents,  like  hydrogen  sulphide,  H2S,  sulphur  dioxide,  SOj,  and  hydri- 
odic  acid,  HI,  reduce  it  to  iodine  : 

HIO3  + 5HI  - 3H3O  + 3lr 

Iodic  Anhydride,  IjO^,  is  a white  erystalline  powder,  which  dis- 
solves in  water  to  form  iodic  acid.  It  decomposes  at  300°  into  iodine 
and  oxygen.  It  can  be  obtained  directly  from  ozone  and  iodine. 

Periodic  Acid,  HIO^.  Normal  periodic  acid  is  not  known.  The 
hydrate,  HIO^ . 2H2O,  is  produced  by  the  action  of  iodine  upon  per- 
chloric acid  : 

2HC10^  -f  I2  + 4H2O  = 2 [HIO^.  2H2O]  + Cb- 

Upon  the  evaporation  of  the  aqueous  solution,  the  acid  crystallizes  in 
colorless  forms,  which  deliquesce,  melt  at  130°,  and  at  about  140°  decom- 
pose into  water,  oxygen  and  periodic  anhydride : 

2(HI0,.  2H2O)  = 1,0,  + O2  + 5H2O. 

Periodates  result  on  conducting  chlorine  into  hot  alkaline  solutions  of 
iodides  or  iodates.  A sodium  periodate  has  been  found  in  Chile  salt- 
peter. 


The  existence  of  the  hydrates  of  periodic  and  perchloric  acids  as  well  as  of  many 
others  (see  Sulphuric  and  Nitric  Acids),  which  were  once  regarded  as  molecular  compounds, 
is  interpreted  at  present  by  the  acceptance  of  hydroxyl  groups,  directly  combined  with 
the  element  of  higher  valence  : 

VII 

CIO^H  -|-  H2O  = €102(011)3,  Trihydrate  or  trihydric  acid. 

VII 

CIO4H  2H2O  = CIO  (OH)^,  Pentahydrate  or  pentahydric  acid. 

VII 

CIO^H  3H2O  = C1(0H)7,  Heptahydrate  or  heptahydric  acid. 

The  maximum  hydrates,  ClfOH).^  and  I(OH).j,  in  which  all  seven  affinities  of  the 
halogen  atom  are  attached  to  hydroxyl  groups,  are  not  known,  but  probably  exist  in 
aqueous  solution.  As  they  give  up  water,  and  one  atom  of  oxygen  becomes  simultaneously 
united  with  two  bonds  to  the  halogen,  they  yield  the  lower  hydrates — even  to  the  mono- 
hydrate CIO3 . OH.  Perchloric  acid  continues  monobasic  in  the  polyhydrates,  since  but  on^ 
hydrogen  atom  is  replaced  by  metals  : 

C\0,U,  -b  KOH  =:  CIO4K  + 3H2O. 

On  the  other  hand,  periodic  acid  (IO3 . OH)  is  not  only  monobasic,  but  as  a pentahydrate, 
I0(0H)5  = HIO4  . 2H2O,  can,  like  the  polybasic  acids,  furnish  also  polymetallic 
salts,  as  : 

VII  / (011)3  VII  f (0H)3  VII  VII 

10  \ (0Na)2  10  \ (0Ag)2  I0(0.Na)5  I0(0Ag)3. 

Salts  also  exist  which  are  derived  from  condensed  polyiodic  acids,  as  : 

^O  , Diperiodic  acid,  etc. 

io<(on)/ 

(Compare  Uisulphuric  Acid,  Dichromic  Acid,  Pyrophosphoric  Acid,  etc.) 


i82 


INORGANIC  CHEMISTRY. 


The  existence  of  such  salts  plainly  indicates  that  the  hydrates  of  acids  must  be  looked 
upon  as  hydroxyl  compounds,  and  that  iodine  and  the  halogens  are,  in  fact,  heptads  in 
their  highest  combinations. 


The  oxygen  compounds  of  the  halogens  in  some  respects  display  a charac- 
ter exactly  opposite  to  that  of  the  hydrogen  derivatives.  While  the  affinity 
of  the  halogens  for  hydrogen  diminishes  with  increasingatomic  weight  from 
fluorine  to  iodine  (see  p.  65),  the  affinity  for  oxygen  is  the  exact  reverse. 
Fluorine  is  not  capable  of  combining  with  oxygen;  the  chlorine  and 
bromine  compounds  are  very  unstable,  and  are  generally  not  known  in 
free  condition  ; the  iodine  derivatives,  on  the  contrary,  are  the  most  stable. 
In  accord  with  this  is  the  fact  that  in  the  higher  oxygen  compounds 
chlorine  and  bromine  are  set  free  by  iodine,  while  in  the  hydrogen  and 
metallic  compounds  of  the  halogens  the  direct  reverse  is  the  case,  viz., 
that  iodine  and  bromine  are  replaced  by  chlorine. 

Further,  the  oxygen  compounds  exhibit  the  remarkable  peculiarity  that 
their  stability  increases  with  the  addition  of  oxygen.  The  lowest  acids, 
HCIO,  HBrO,  are  very  unstable,  even  in  their  salts;  they  possess  a very 
slight  acid  character,  and  are  separated  from  their  salts  by  carbon 
dioxide.  The  most  energetic  and  most  stable  are  the  highest  acids, 
HCIO4,  HBrOg,  HIO^,  in  which  the  higher  valence  of  the  halogens 
appears.  In  the  sulphur  and  nitrogen  groups  those  oxides,  in  which  the 
elements  manifest  their  maximum  valence,  are  the  most  stable  (compare 
p.  169). 


2.  OXYGEN  COMPOUNDS  OF  THE  ELEMENTS  OF 
THE  SULPHUR  GROUP. 


The  elements  sulphur,  selenium,  and  tellurium  combine  with  oxygen  in 
several  proportions.  Sulphur  and  oxygen  form  the  following  derivatives : 


S,03 

Sulphur 

sesquioxide. 

SO3 

Sulphur  trioxide, 
Sulphuric  anhydride. 


SO2 

Sulphur  dioxide, 
Sulphurous  anhydride. 

S2O,. 

Sulphur  heptoxide, 
Persulphuric  anhydride. 


The  oxygen  derivatives  of  selenium  and  tellurium  correspond  to  sulphur 
di-  and  trioxides : 

Se02  wSeOj 

Selenium  dioxide,  Selenic 

Selenious  anhydride.  anhydride. 


TeOj 

Tellurous 

anhydride. 


TeOg. 

Telluric 

anhydride. 


Each  of  these  oxides,  sulphur  sesquioxide  excepted,  combines  with  one 


OXYGEN  COMPOUNDS  OF  SULPHUR.  1 83 

molecule  of  water  to  form  a dibasic  acid  (p.  172),  of  which  the  respective 
oxide  is  the  anhydride.  The  acids  of  sulphur  are : 

H2SO3  HjSO^  H2S20g. 

Sulphurous  Sulphuric  Persulphuric 

acid.  acid.  acid. 

By  the  replacement  of  one  atom  of  hydrogen  by  one  atom  of  metal  the 
so-called  acid  or  primary  salts  result,  while  the  neutral  or  secondary  salts 
are  obtained  by  the  replacement  of  both  hydrogen  atoms : 

KHSO^  K2SO,. 

Acid  potassium  sulphate,  Neutral  potassium  sulphate, 

Monopotassium  sulphate.  Dipotassium  sulphate. 


1.  OXYGEN  COMPOUNDS  OF  SULPHUR. 

Sulphur  eombines  in  four  proportions  with  oxygen.  There  are  in  addi- 
tion nine  compounds  which  contain  hydrogen  besides  these  two  elements ; 
they  are  dibasic  acids.  The  only  one  which  can  be  readily  prepared  is 
sulphuric  acid.  The  others  are  stable  only  in  aqueous  solution  or  in  the 
form  of  salts.  In  the  table  which  follows  the  anhydrides  are  arranged 
opposite  to  the  acids  into  which  they  pass  by  the  addition  of  water. 


Oxides— Anhydrides. 


Sulphur  sesquioxide,  SjOg 
Sulphur  dioxide,  SOj 
Sulphur  trioxide,  SO, 

Sulphur  heptoxide,  SgO^ 

Polythionic 

acids 


Both  sulphur  dioxide  and  trioxide 
di-acid  ox  pyro-acid : 

2502  + H2O  = H2S2O5,  Disulphurous  acid. 

2503  -j-  HjO  = H2S2O7,  Disulphuric  acid. 

The  latter  alone  is  known  in  a free  state. 

Sulphur  Dioxide,  SO2,  or  sulphurous  anhydride,  is  formed  by  burn- 
ing sulphur  or  sulphides  in  the  air  : 

S -I  O2  = SO2. 

I vol.  I vol. 

A little  trioxide  is  always  produced. 

It  may  also  be  obtained  by  heating  sulphur  with  the  oxides  of  copper, 
manganese  and  lead : 

2CuO  -f  2S  = CujS  + SOj. 

Cuprous 

sulphide. 


Acids. 

Thiosulphuric  acid,  HjSjOj 

Hyposulphurous  acid,  HjSOg 

Sulphurous  acid,  HjSO, 

Sulphuric  acid,  HjSO^ 

Persulphuric  acid,  H2S2O8 

Dithionic  acid,  HgSjOg 

(Hyposulphuric  acid) 
Trithionic  acid,  HjSgOg 

Tetrathionic  acid,  H2S4O5 

Pentathionic  acid,  H,ScOc. 


yield  a so-called  anhydro-acid^ 


i84 


INORGANIC  CHEMISTRY. 


It  is  most  conveniently  prepared  for  laboratories  by  heating  concentrated 
sulphuric  acid  (i  part)  with  coi)per  jiart)  : 

2H2SO4  + Cu  = CuSO^  -h  SO2  4-  2H2O. 

Copper 

sulphate. 

Usually  a little  cuprous  sulphide  separates  : 

5Cu  + 411280,  = CU2S  + 3CUSO,  -f  41130. 

This  is  due  to  a far-reaching  reduction. 

Sulphuric  acid  is  similarly  decomposed  (reduced)  by  heating  it  with 
carbon  : 

2lbSO,  -f  C = 2SO2  4-  CO2  + 2II2O. 

By  this  method  we  get  a mixture  of  carbon  and  sulphur  dioxides,  which 
are  separated  with  difficulty.  A more  convenient  method  for  prepar- 
ing sulphur  dioxide  consists  in  allowing  ordinary  sulphuric  acid  to 
act  upon  calcium  sulphite,  CaSOg.  The  latter  is  mixed  with  burnt 
gypsum  {yz  part)  and  water,  then  moulded  into  cubes,  which  can  be 
introduced  into  a Kipp  generator,  as  in  the  preparation  of  oxygen 
(p.  81).  Owing  to  its  solubility  in  water,  sulphur  dioxide  must  be  col- 
lected over  mercury. 

Sulphur  dioxide  is  a colorless  gas,  with  a suffocating  odor.  One  liter 
of  it  weighs  2.8615  grams  under  normal  conditions.  Its  specific  gravity 
equals  2.21  (air=  i)  or  64.06  (O2  = 32),  corresponding  to  the  molecular 
formula  SOj.  It  condenses  at  — 15°,  or  at  ordinary  temperatures  under 
a pressure  of  three  atmospheres,  to  a colorless  liquid,  of  specific  gravity 
1.43  at  0°,  which  crystallizes  at  — 76°  and  boils  at  — 8°.  Its  critical 
temperature  is  157°;  its  critical  pressure  79  atmospheres.  Upon  evapo- 
ration the  liquid  sulphur  dioxide  absorbs  much  heat  ] being  easily 
accessible  it  is  used  in  ice  machines.  If  some  of  the  liquid  is  poured 
upon  mercury  in  a clay  crucible,  and  the  evaporation  accelerated  by 
blowing  air  upon  it,  the  metal  will  solidify.  Water  dissolves  50  volumes 
of  sulphur  dioxide  gas  with  liberation  of  heat.  The  gas  is  again  set  free 
upon  application  of  heat.  The  solution  shows  all  the  chemical  properties 
of  the  free  gas. 

Sulphur  dioxide  has  great  affinity  for  oxygen.  The  gases  combine 
when  dry,  if  their  mixture  be  conducted  over  feebly  heated  platinum 
sponge ; * sulphur  trioxide  results : 

2SO2  + O.,  = 2SO3. 

2VOls.  IVOl. 

Winkler’s  method  for  producing  sulphuric  acid  technically  is  based  on 
this  reaction  (compare  p.  193). 

In  a(pieous  solution  the  dioxide  slowly  absorbs  oxygen  from  the  air,  and 
becomes  sulphuric  acid  : 

S()2  + H20-f  0=11.280,. 


* Instead  of  platinum  s])onge,  platinized  asbestos  may  be  applied  ; this  is  obtained  by 
immersin}:^  asbestos  in  a platinic  chloride  solution,  then  in  ammonium  chloride,  and  after- 
ward drying  and  igniting. 


SULPHUROUS  ACID.  185 

Aqueous  sulphur  dioxide  is  converted  more  rapidly  into  sulphuric  acid 
by  the  action  of  the  halogens  chlorine,  bromine  and  iodine: 

H,S03  + H^O  -f  Cl,  = H,SO,  + 2HCI. 

Here  the  decomposition  of  a molecule  of  water  is  effected  in  conse- 
quence of  the  affinity  of  the  halogen  for  hydrogen  and  of  sul})hurous  acid 
for  oxygen.  On  adding  sulphurous  acid  to  a dark-colored  iodine  solution 
the  latter  is  decolorized. 

Similarly,  sulphurous  anhydride  and  its  solution  withdraw  oxygen  from 
many  compounds  rich  in  that  element ; hence  it  deoxidizes  strongly  and 
passes  over  into  sulphuric  acid.  Thus  chromic  acid  is  reduced  to  oxide, 
and  the  red  solution  of  permanganic  acid  is  decolorized  with  formation 
of  manganous  salts.  Many  organic  coloring  substances,  like  those  of 
flowers,  are  decolorized  by  it.*  This  property  is  what  leads  to  its  appli- 
cation in  the  bleaching  of  wools  and  silks,  which  are  strongly  attacked 
by  the  ordinary  chlorine  bleaching  agents  (p.  53). 

Again,  the  dioxide  may  be  deoxidized  by  stronger  reducing  agents  (it 
acts  with  them  as  an  oxidant);  thus  in  the  presence  of  water,  sulphur  is 
separated  from  it  by  hydrogen  sulphide  : 

SO2  + 2H2S  = 2H2O  + 3S. 

If,  however,  both  gases  are  perfectly  dry  or  strongly  diluted  by  other 
neutral  gases,  the  action  is  very  slow.  (See  Hydrosulphurous  and  Penta- 
thionic  Acids.) 


A mixture  of  equal  volumes  of  sulphur  dioxide  and  chlorine  unites  in  direct  sunlight 
to  sulphuryl  chloride  SO^Ch  (p.  195).  When  sulphur  dioxide  acts  upon  warmed 
phosphoric  chloride,  the  products  are  phosphorus  oxychloride,  and  the  compound  SOC’b  : 

SO2  + PCI5  = POCI3  -f  SOCI2. 

Chlorthionyl,  SOC^,  may  be  viewed  as  the  chloride  of  sulphurous  acid  or  as  sul- 
phur dioxide  in  which  one  atom  of  oxygen  is  replaced  by  two  atoms  of  chlorine  (p.  186). 
Thionyl  chloride  may  be  made  by  the  interaction  of  sulphur  dichloride  and  sulphur  tri- 
oxide : 

SO3  + SCI2  = SOCI2  + SO2 

[see  Michaelis,  Ann.  Chem.  274  (1894),  184].  It  is  a colorless  liquid  with  a sharp 
odor,  and  boils  at  78°.  Its  specific  gravity  equals  1.67.  Water  decomposes  it  into 
hydrogen  chloride  and  sulphurous  acid  : 

SOCI2  + H2O  = SO2  + 2PICI. 

Sulphurous  Acid,  H2SO3,  is  not  known  in  free  condition,  but  is 
probably  present  in  the  aqueous  solution  of  sulj^hur  dioxide.  On  cool- 
ing the  concentrated  solution  to  0°,  colorless  cubical  crystals  separate, 
containing  probably  six  molecules  of  water  (Geuther).  If  the  aqueous 
solution  is  allowed  to  stand  for  some  time,  especially  in  sunlight,  sulphur 
separates  with  the  formation  of  sulphuric  acid  : 

3SO2  -p  2ii20  = ri22SO^  -j-  s. 


*The  acid  forms  colorless  compounds  with  dyestuffs  of  the  flowers.  Dilute  sulphuric 
acid  or  heat  breaks  down  these  derivatives,  i.  e.,  the  original  colors  reappear. 


i86 


INORGANIC  CHEMISTRY. 


Sulphurous  acid  is  dibasic  and  forms  two  series  of  salts;  the  jirimary 
(KHSO3)  and  secondary  (KjSOg). 

The  sulphites,  with  the  exception  of  those  of  the  alkalies,  are  insolu- 
ble or  dissolve  with  difficulty  in  water  [see  Seubert  and  Elten,  Z,  f. 
anorg.  Ch.  4 (1893),  44].  When  sulphurous  acid  is  separated  out 
from  its  salts  by  stronger  acids  it  decomposes  into  its  anhydride  and  water : 

Na^SOg  -f  2IICI  = 2NaCl  -f  SO,  + II^O. 


The  following  is  all  that  is  known  regarding  the  chemical  structure  of  sulphurous  acid 
and  its  derivatives.  Its  anhydride  and  chloride  have  the  formulas  : 

IV  IV 

O = S = O and  O = 

By  water  absorption  or  by  decomposition  with  water  the  hydrate 


IV 

O = S< 


OH 

OH 


results.  This  formula  indicates  that  the  hydrogen  atoms  or  the  metals  which  may  replace 
them  are  not  directly  combined  with  the  sulphur,  but  are  linked  through  oxygen.  Or- 
ganic compounds,  esters  of  sulphurous  acid,  are  known  which  undoubtedly  are  derived 
from  this  symmetrical  formula.  The  inorganic  salts  of  the  acid,  however,  very  probably 
contain  one  metal  atom  in  direct  union  with  sulphur,  hence  their  basal  acid  must  have 
the  formula 


IV 

O = S< 


O-OH 

H 


O^vi  OH, 


accordingly  as  sulphur  is  regarded  as  quadrivalent  or  sexivalent.  Organic  derivatives  are 
also  known  of  this  unsymmetrical  acid.  From  this  it  would  appear  that  the  anhydride 
readily  yields  compounds  which  may  be  derived  either  from  a symmetrical  or  an  un- 
symmetrical hydrate,  SOgHj.  See  Sodium  Sulphite  and  also  Richter’s  Organic  Chemistry. 

The7’ino-che7nical  Depo7-t77ie7it. — Sulphur  dioxide  is  a very  powerful  exothermic  com- 
pound. 71. 1 Cal.  are  set  free  in  its  formation  from  rhombic  sulphur  and  oxygen.  When 
it  dissolves  in  much  water  there  is  an  additional  disengagement  of  7.7  Cal.,  so  that  the 
heat  of  formation  of  the  hypothetical  sulphurous  acid  in  dilute  aqueous  solution  (from 
sulphur,  oxygen  and  water)  equals  78.8  Cal.  : 


(S,02)  gas  = 7I-I  ; (S02,Aq)  = 7.7  ; (S,02,Aq)  = 78.8. 


In  consequence  of  this  great  loss  of  energy  the  dioxide  is  a very  stable  compound ; it 
is  only  at  high  temperatures  that  it  sustains  a partial  separation  into  sulphur  and  oxygen. 
For  its  behavior  toward  oxygen,  see  p.  187. 


Hydrosulphurous  Acid,  H.2SO2  or  H2S2O4.  On  adding  zinc,  iron,  and  some  other 
metals  to  the  aqueous  solution  of  sulphurous  acid  they  dissolve  without  liberation  of 
hydrogen  to  yellow-colored  liquids.  Schonbein  (1852)  observed  that  such  solutions  pos- 
sessed, when  ajjplied  to  indigo,  strong  bleaching  properties,  and  assumed  that  there  was 
in  them  a jjeculiar  acid,  the  composition  of  which  was  first  determined  by  Schiitzenberger 
in  1869.  The  hydrosulphurous  acid  is  formed  there  by  the  action  of  the  hydrogen  set 
free  by  the  zinc  upon  a second  molecule  of  sulphurous  acid  : 

H2SO3  + Zn  = ZnSOg  + H2 

and 

H2SO3  + H2  = H2SO2  + FI2O. 

The  aqueous  solution  of  the  acid  has  an  orange-yellow  color,  reduces  powerfully,  bleaches 
and  soon  decomposes  with  separation  of  sulphur  and  the  formation  of  sulphurous  acid. 

'I'lie  salts  are  more  stable  than  the  acid.  The  sodium  salt  is  obtained  by  the  action  of 
zinc  filings  upon  a concentrated  solution  of  primary  sodium  sulphite.  It  is  used  in  dye- 


SULPHUR  TRIOXIDE. 


187 


ing  and  cotton  printing  to  bleach  indigo.  Its  composition  is  not  established  with  cer- 
tainty ; it  corresponds  to  either  the  formula  NaHSOg  or  NujSjO^.  The  salt  solutions 
absorb  oxygen  very  rapidly  from  the  air  and  change  to  sulphites. 


Sulphur  Sesquioxide,  SjOg,  is  obtained  by  the  solution  of  flowers  of  sulphur  in 
liquid  sulphuric  anhydride  ; it  separates  out  in  blue  drops,  which  solidify  to  a mass  resem- 
bling malachite.  It  decomposes  gradually,  more  rapidly  on  warming,  into  sulphur  dioxide 
and  sulphur.  It  is  very  violently  broken  down  by  water,  with  formation  of  sulphur, 
sulphurous,  sulphuric,  and  polythionic  acids.  It  dissolves  with  a blue  color  in  fuming 
sulphuric  acid. 

Sulphur  Trioxide,  SO3,  or  sulphuric  anhydride,  is  produced,  as  pre- 
viously described,  by  the  union  of  sulphur  dioxide  and  oxygen,  aided  by 
heated  platinum  sponge.  Platinized  balls  of  white  clay  are  used  in  tech- 
nical operations.  It  is  also  formed  when  sulphur  dioxide  and  air  are  con- 
ducted over  ignited  oxide  of  iron,  chromic  oxide,  or  manganese  oxide. 
These  oxides  act  like  the  platinum  sponge,  platinized  asbestos  or  clay. 
They  are  merely  contact  substances.  It  can  also  be  made  by  heating 
sodium  or  potassium  pyrosulphate  (p.  194)  and  anhydrous  sulphates,  e.  g., 
ferric  sulphate : 

FcjlSOJa  = Fe^Oa  + 3SO3  ; 

and  is  most  conveniently  obtained  by  heating  fuming  (Nordhausen)  sul- 
phuric acid  (p.  193);  the  escaping  white  fumes  are  condensed  in  a chilled 
receiver.  It  may  be  obtained  pure  by  repeated  distillation,  by  fusion  at 
moderate  temperature  (20-30°),  and  then  pouring  it  off  from  the  re- 
maining solid  portions.  It  crystallizes  in  long,  broad,  transparent 
needles,  which  melt  at  14.8°  to  a very  mobile  liquid.  At  16°  its  specific 
gravity  is  1.940.  It  distils  at  46°.  The  vapor  density  agrees  with  the 
formula  SO3.  The  perfectly  pure  anhydride  does  not  change  on  preser- 
vation. If,  however,  by  absorption  of  water  it  contains  traces  of  sulphuric 
acid,  it  soon  becomes  an  asbestos-like  mass,  which  does  not  melt  until  at 
about  50°.  This  was  formerly  regarded  as  a peculiar  form  of  sulphuric 
anhydride.  The  pure  anhydride  can  be  readily  obtained  from  it  by 
distillation  [R.  Weber;  see  also  Rebs,  Ann.  Chem.  246  (1888),  379]. 

Sulphuric  oxide  fumes  strongly  in  the  air,  and  attracts  moisture  with 
avidity.  When  thrown  on  water  it  dissolves  with  hissing  to  form  sul- 
phuric acid  (SO3  -f-  H2O  = H2SOJ. 

When  the  vapors  are  led  through  heated  tubes  they  are  decomposed 
into  sulphur  dioxide  and  oxygen. 


Thermo- chemical  Deportment. — When  sulphur  dioxide  and  oxygen  combine  to  form 
liquid  sulphur  trioxide  32. 1 Cal.  are  disengaged,  so  that  its  heat  of  formation  from  the 
elements  is  103.2  Cal.  : 

(S0^,0)  liquid  = 32.1,  (8,03)  liquid  = 103.2, 

inasmuch  as  the  heat  of  formation  of  the  dioxide  71.  i Cal.  (p.  186).  This  is  another 
contradiction  of  Berthelot’s  princi{)le  of  the  greatest  evolution  of  heat.  According  to  it 
when  sulphur  burns  in  the  air  or  in  oxygen  it  should  form  not  the  dioxide  but  the  trioxide, 
because  in  the  latter  case  there  would  occur  the  greater  heat  evolution.  The  fact  is  that 


i88 


INORGANIC  CHEMISTRY. 


from  91. 1-95.5  percent,  of  the  sulphur  is  converted  into  dioxide,  and  only  2. 9-2. 5 per 
cent,  into  trioxide.  In  the  presence  of  porous  bodies  (ferric  oxide,  etc.)  the  <)uantity 
of  trioxide  reaches  as  much  as  13  per  cent. 

On  dissolving  .sulphur  trioxide  in  much  water  to  form  afjueous  sulphuric  acid,  39.2 
Cal.  are  disengaged.  The  j)roduction,  therefore,  of  the  aqueous  acid  from  sulphur, 
oxygen,  and  water  equals  (including  the  heat  of  formation  of  SO3)  142.4  Cal.  : 

(S03,Aq)  = 39.2  ; (S,03,Aq)  = 142.4. 

If  we  add  to  this  the  heat  of  formation  of  water  (liquid),  68.3  Cal.,  the  heat  of  f«^r 
mation  of  sulphuric  acid  (II2SO4  — SO3  -(-  HjO)  from  the  elements  in  dilute  aqueous 
solution  will  be  : 

(Il2,S,04,Aq)  = 210.7. 

The  heat  of  .solution  of  anhydrous  sulphuric  acid,  in  much  water,  equals  17.8  Cal.  ; 
hence  the  heat  of  formation  of  anhydrous  sulphuric  acid  from  its  elements  is  210.7  — 
17.8  = 192.9  : 

(Il2S04,Aq)  = 17.8  ; (Il2,S,04)  = 192.9. 

Sulphur  Heptoxide,  was  obtained  by  Berthelot  on  conducting  a silent  elec- 

tric discharge  of  considerable  intensity  through  a mixture  of  equal  volumes  of  dry  .sulphur 
dioxide  and  oxygen.  It  separates  in  oily  drops  which  .solidify  at  0°  to  a crystalline  mass. 
This  compound  must  be  regarded  as  the  anhydride  of  persulphuric  acid.  It  decomposes 
upon  standing,  immediately  when  heated,  into  oxygen  and  sulphur  trioxide  : 

S2O7  = 2SO3  + O. 

It  fumes  strongly  in  the  air  and  with  water,  as  with  heat,  decomposes  into  sulphuric 
acid  and  oxygen  : 

^2^7  "f"  2H2O  2H2SO4  -J-  O. 

Persulphuric  Acid,  H2S20g,  corresponding  to  sulphuric  heptoxide,  has  not  been 
obtained  in  a pure  condition.  A solution  of  it  may  be  formed  at  the  anode  when  sul- 
phuric acid  (cooled,  40  per  cent.)  is  electrolyzed.  Richarz  thinks  that  it  is  produced  in 
this  instance  by  the  sulphuric  acid  breaking  down  into  the  ions  H and  HSO4  (p.  92), 
the  latter  then  combining  to  1^28203.  Inactive  oxygen  and  ozone  are  formed  simulta- 
neously. A solution  of  the  acid  can  be  obtained  by  dissolving  sulphur  heptoxide  in  dilute 
sulphuric  acid,  also  by  the  addition  of  aqueous  hydrogen  peroxide  to  cooled,  concentrated 
sulphuric  acid,  when  ozone  is  also  produced. 

The  solution  of  persulphuric  acid  in  sulphuric  acid  exhibits  oxidation  reactions  similar 
to  those  of  hydrogen  peroxide.  It  oxidizes  ferrous  to  ferric  sulphate,  gradually  liberates 
iodiive  from  potassium  iodide  and  decolorizes  the  blue  solution  of  indigo-sulphuric  acid  ; 
however,  it  is  not  capable  of  decolorizing  a permanganate  solution,  neither  does  it  oxi- 
dize chromic  to  perchromic  acid  or  affect  titanic  acid  solutions  (p.  103). 

Hydrogen  peroxide  is  produced  if  the  electrolyzed  sulphuric  acid  contains  more  than 
60  per  cent,  of  sulphuric  acid  ; this  is  caused  by  the  breaking  down  of  the  persulphuric 
acid  which  has  been  formed  (p.  loi). 

d'he  persulphates  result  in  the  action  of  strong  bases  upon  the  heptoxide,  or,  as  shown 
by  Hugh  Marshall  and  Berthelot,  by  the  electrolysis  of  sulphate  solutions  [Elbs,  Jr. 
j)rakt.  Ch.  48  (1893),  185,  and  Chem.  Zeit.  1895,  1120].  They  will  be  described  under 
the  respective  metals.  Their  dilute  solutions  show  the  following  properties,  character- 
istic for  persulphuric  acid  : when  heated  they  give  out  ozone,  and  in  the  presence  of 
hydrochloric  acid  chlorine  ; they  precipitate  the  hydrated  dioxide  of  manganese  upon 
the  addition  of  a manganous  salt ; in  the  presence  of  sulphuric  acid  they  oxidize  aniline 
to  aniline  black.  See  p.  198  for  the  probable  structure  of  persulphuric  acid. 


SULPHURIC  ACID. 


189 


SULPHURIC  ACID. 

H2SO4. 

it  is  certain  that  this  acid  was  known  in  the  fifteenth  century,  probably 
long  before  that  time.  It  is  the  most  important  of  all  the  acids;  nearly 
all  of  them  can  be  prepared  directly  or  indirectly  by  means  of  it.  In 
technical  chemistry  and  in  the  arts  it  meets  with  an  unusually  exten- 
sive and  varied  application,  but  it  is  especially  valuable,  in  the  Le  Blanc 
soda  process,  in  the  refining  of  petroleum  and  tar-oils,  for  the  production 
of  aluminium  sulphate,  blasting  material,  dyes,  artificial  manures,  etc., 
etc. 

Besides  the  reactions  already  mentioned,  it  is  formed  in  the  oxidation  of 
sulphur  by  nitric  acid.  It  was  obtained  formerly  by  heating  ferrous  sul- 
phate (FeSOJ,  which  breaks  down  into  ferric  oxide,  sulphurous  acid  and 
sulphuric  anhydride : 

2FeSO^  = FejOj  SOj  + SO3. 

At  present,  however,  it  is  almost  exclusively  manufactured  in  large  quan- 
tities, after  the  so-called  English  lead-chamber  process.  This  method  is 
based  upon  the  conversion  of  sulphur  dioxide  into  sulphuric  acid  by  means 
of  nitric  acid.  Sulphur,  pyrite  (FeSj),  or  other  blendes  are  roasted  in 
ovens,  and  the  disengaged  sulphur  dioxide  immediately  conducted,  to- 
gether with  air,  into  a series  of  large  leaden  chambers  in  which  it  is  fre- 
quently brought  in  contact  with  nitric  acid  and  steam. 

The  nitric  acid  gives  up  a portion  of  its  oxygen  to  the  sulphur  dioxide 
and  thereby  oxidizes  it  in  the  presence  of  water  to  sulphuric  acid,  which 
collects  on  the  floors  of  the  lead  chambers.  The  nitrogen  oxides,  NO, 
NOj,  and  N2O3,  arising  simultaneously  from  the  nitric  acid,  are  capable 
of  transferring,  in  various  ways,  the  oxygen  of  the  air,  which  enters  the 
chambers,  in  the  presence  of  water,  to  the  sulphurous  acid  so  that  it 
passes  over  into  sulphuric  acid.  As  the  nitrogen  oxides  are  being  con- 
stantly regenerated,  a given  quantity  of  nitric  acid  should,  according  to 
theory,  be  capable  of  converting  unlimited  amounts  of  sulphur  dioxide 
into  sulphuric  acid  if  the  water  and  oxygen  are  present  in  sufficient 
quantity.  Facts  show  that  this  is  really  not  true,  because  a portion  of 
the  nitric  acid  (1-2  parts  by  weight  for  100  parts  of  sulphurous  acid)  is 
reduced  to  nitrous  oxide,  nitrogen,  and  probably  ammonia, — compounds 
which  do  not  participate  in  the  oxidation  of  the  sulphurous  acid.  Hence, 
in  the  lead  chambers  the  nitrogen  oxides,  particularly  nitric  oxide,  play 
the  role  of  oxygen  carriers. 

The  chemical  changes  occurring  side  by  side  in  the  lead  chambers  are  influenced  by 
the  quantities  of  the  reacting  substances  and  by  the  temperature  ; consequently,  they  are 
not  the  same  for  all  chambers,  indeed  not  the  same  for  different  parts  of  the  same  cham- 
ber. The  chemical  and  physical  conditions  are  being  constantly  altered  by  the  violent 
mixing  together  of  the  gases,  and  it  seems  almost  a fruitless  task  to  attempt  to  establish  a 
definite  theory  for  the  sul])huric  acid  manufacture.  The  most  important  of  these  changes 
appear  to  be  the  following  : 

In  the  presence  of  water,  the  nitric  acid  oxidizes  the  sulphur  dioxide  to  sulphuric  acid, 
and  the  former  is  reduced  to  nitric  oxide  (NO)  or  nitrogen  dioxide  (NOj)  : 

3SO2  -f  2IINO3  -f  2II2O  = 3H.,S0,  1 2NO. 


190 


INORGANIC  CHEMISTRY. 


The  nitric  oxide  unites  with  the  oxygen  of  the  air  (which  entered  the  chambers  simul- 
taneously with  the  sulphur  dioxide)  to  form  nitrogen  dioxide,  which,  in  the  presence  of 
water,  converts  a fresh  portion  of  sulphur  dioxide  into  sulphuric  acid  : 

SO2  + 11,0  + NO,  = HjSO*  + NO. 

Tlie  regenerated  nitric  oxide  is  again  subjected  to  the  same  transformations.  Tn  Tunge’s 
ojunion  the  lead-chamber  process  is  not  an  alternating  reduction  and  oxidation  of  the  oxides 
of  nitrogen,  but  rather  a condensation  of  nitric  oxide  and  nitrogen  dioxide  with  sulphur 

dioxide,  oxygen,  and  water  to  nitrosylsulphuric  acid,  (P-  207)  : 

2SO,  + 2NO  O3  -t- 11,0  = 2S0,<0-,N0 

The  excess  of  water  in  the  lead  chamber  immediately  converts  this  product  into  sulphuric 
acid  and  nitrogen  trioxide  or  a mixture  of  nitric  oxide  and  nitrogen  dioxide  : 

2SO,<2i- ■+  N,0.. 

N,03  = NO  + NO, ; N,03  II,0  ^ 2IINO,. 

In  the  anterior  portion  of  the  chamber  system  the  nitrosylsulphuric  acid  breaks  down 
with  the  assistance  of  the  sulphurous  acid  and  with  the  formation  of  nitric  oxide  : 

2S0,<0jj^°  + so,  + 211,0  = 3S0,(0H),  + 2NO. 

The  regenerated  nitric  oxide  and  nitrogen  dioxide  act  together  with  air  and  steam  upon 
new  quantities  of  sulphur  dioxide  in  the  manner  already  indicated.  According  to  Lunge, 
nitrosylsulphuric  acid  is  formed  upon  the  very  first  action  of  the  nitric  acid  introduced  into 
the  lead  chamber  : 

SO,  NO, . OH  = 

and  it  then  reacts  in  the  manner  indicated.  [See  Lunge,  Z.  f.  anorg.  Ch.  7 (1894), 
212  ; R.  Hasenclever,  Ber.  29  (1896),  iii,  2861  ; H.  Ost,  Lehrbuch  des  technischen 
Chemie,  3 Aufl.  (1898),  and  especially  Jurisch,  in  Dammer’s  Handbuch  der  chemischen 
Technologic  (1895)  I,  163.] 

In  the  lead-chamber  process  the  active  nitrogen  oxides  (NO  and  NO,)  are  carried  along 
and  withdrawn  from  the  action  by  means  of  the  escaping  nitrogen  and  excess  of  air.  To 
avoid  any  further  loss  of  nitric  acid  by  this  means,  the  escaping  brown  gases  are  conducted 
through  the  so-called  Gay-Lussac  tower.  This  is  constructed  of  lead  sheets,  and  filled 
with  pieces  of  coke,  or,  as  these  are  apt  to  become  coated  with  a mud  of  lead  sulphate 
which  stops  them  up,  with  fire-brick  or  cylinders  over  which  concentrated  sulphuric  acid 
constantly  trickles.  The  acid  completely  absorbs  the  nitrogen  oxides  K^O,,  NO,  and 
NO,  with  formation  of  nitrosylsulphuric  acid  (see  p.  208).  The  nitrogen  oxides  can  be 
regained  from  the  acid — the  so-called  nitroso-acids — collected  at  the  bottom  of  the  tower, 
and  made  useful  in  the  production  of  sulphuric  acid  in  the  chambers.  This  is  effected  at 
present  in  the  so-called  Glove)'  to^veVy  which  is  constructed  of  lead  plates  and  fire-proof 
bricks,  and  inserted  between  the  sulphur  ovens  and  lead  chambers.  In  this  the  nitroso- 
acid  (diluted  with  the  previously  obtained  chamber  acid)  is  allowed  to  run  over  fire-brick, 
while  the  hot  gases  of  combustion  from  the  sulphur  ovens  stream  against  it.  This  cools 
the  hot  gases  to  the  required  temperature  (70-80°),  water  evaporates  from  the  chamber 
acid,  and,  at  the  same  time,  the  nitrogen  oxides  are  set  free  (see  above),  and  carried 
into  the  lead  chambers.  Hence,  the  (Hover  tower  serves  not  only  for  complete  utiliza- 
tion of  the  nitrogen  oxides,  but  also  for  the  concentration  of  the  chamber  acid  (to  82 
j)er  cent.  I I,SOJ. 

'I'he  chamber  j)rocess  may  be  illustrated  by  the  following  experiment : A large  glass 
flask  A (Fig.  64)  replaces  the  lead  chamber;  in  its  neck  are  introduced,  by  means  of  a 
cork,  several  glass  tubes,  which  serve  to  introduce  the  various  gases.  In  rr,  sulphur  di- 


SULPHURIC  ACID. 


I9I 


oxide  is  generated  by  heating  a mixture  of  sulplmric  acid  and  mercury  or  copper  turn- 
ings. The  flask  b contains  some  dilute  nitric  acid  and  copper  turnings,  from  wliich  nitric 
oxide  (NO)  is  evolved.  Water  is  boiled  in  c to  get  steam.  Air  enters  through  d while 
the  excess  of  gases  escapes  through  e.  By  the  meeting  of  nitric  oxide  (NO)  with  the 
air,  red  fumes  of  nitrogen  dioxide  (NO2')  arise,  and  these  in  presence  of  water  change 
the  sulphur  dioxide  to  sulphuric  acid.  The  regenerated  nitric  oxide  yields  nitrogen  di- 
oxide with  the  oxygen  of  the  air,  and  converts  another  portion  of  sulphur  dioxide  into 
sulphuric  acid.  In  time  aqueous  sulphuric  acid  collects  upon  the  bottom  of  the  vessel. 

If,  at  first,  only  sulphur  dioxide,  nitric  oxide  and  air  enter  without  the  steam,  we  get 


(by  aid  of  the  moisture  of  the  air)  the  compound  SO2 


{ 


O.NO 

OH 


(the  so-called  ni tro.su Iphuric 


acid)  which  covers  the  walls  of  the  vessel  with  a white  crystalline  sublimate.  These 
crystals,  known  as  lead-chamber  crystahy  are  also  formed  in  the  technical  manufacture 


Fig.  64. 


of  sulphuric  acid,  when  an  insufficient  quantity  of  steam  is  conducted  into  the  chambers. 
Water  decomposes  them  into  sulphuric  acid  and  nitrogen  oxides. 

The  acid  collecting  in  the  chambers  (chamber  acid)  possesses,  when 
the  operation  has  been  properly  conducted,  the  specific  gravity  of  1.5  (50^ 
according  to  Beaiime) ; it  contains  about  60  per  cent,  of  sulphuric  acid 
and  40  ])er  cent,  of  water.*  For  concentration  the  chamber  acid  is  first 
heated  in  lead  jians  until  the  specific  gravity  reaches  1.71  (60°  Beaume ; 


*The  Beaum6  araeometer  {apuLoq,  thin) — hydrometer — used  so  much  in  technical  work 
has  an  empirical  scale.  For  liquids  heavier  than  water  an  instrument  is  used  the  zero 
j)oint  of  which  is  obtained  by  placing  it  in  pure  water.  This  will  give  the  highest  point 
of  the  scale.  A second  fixed  point  is  obtained  by  dipping  the  instrument  into  a solution 
of  15  parts  of  salt  in  85  parts  of  water.  The  distance  between  the  two  points  is  divided 
into  fifteen  equal  parts  or  degrees.  This  division  is  also  continued  further  downward. 


192 


INORGANIC  CHEMISTRY. 


78  per  cent.  H^SOJ.  The  lead  pans  are  strongly  attacked  by  further 
evaporation.  However,  in  vessels  of  cast  iron,  an  acid  of  specific  gravity 
1.8  (64.2°  Beaume  ; 86.9  per  cent.  HjSOj  may  be  olitained.  d'hese 
are  the  crude  sulphuric  acids  (^Acidum  sulphuriciwi  crudutn)  of  commerce. 
They  still  contain  arsenic,  .selenium,  iron  and  lead,  d'he  jiurified  acid 
obtained  in  various  ways  and  diluted  is  again  reduced  in  lead  vessels  to  60° 
Beaume;  further  concentration  to  65.5-66°  Beaume  (92-94  per  cent. 
HjSO^;  specific  gravity  1.83-1.837)  is  conducted  in  vessels  of  glass, 
porcelain  or  platinum.  Tlie  most  concentrated  acid  (98-98.5  percent. 

always  made  in  vessels  of  platinum,  which  in  recent  years  have 
been  lined  within  with  gold,  because  they  are  then  less  attacked. 

By  the  distillation  of  the  crude  bhiglish  acid  an  acpieous  solution  at  first 
distils  over  (one-third  of  the  distillate),  but  at  338°  we  obtain  almost  pure 
sulphuric  acid  {Acidum  sulphuricufu purum  or  destillatuvi).  This  has  the 
specific  gravity  1.854  at  0°  or  1.842  at  12°,  and  contains  about  1.5  per 
cent,  of  water.  On  cooling  this  to  — 35°  white  crystals  separate,  which 
after  repeated  recrystallization  melt  at  10.5°  ; this  is  the  anhydrous  acid, 
HjSO^;  also  called  suljihuric  acid  nionohydrate.  The  crystalline  acid 
is  more  readily  obtained  by  cooling  the  96-98  i)er  cent,  sulphuric  acid 
to  0°  or — 10°,  and  then  adding  already  formed  crystals.  This  is  the 
manner  in  which  the  anhydrous  acid  is  produced  technically  ; the  crys- 
tals are  separated  from  the  liquid  hydrous  acid  by  a centrifugal  machine. 

Pure  anhydrous  sulphuric  acid,  H2SO^  (called  monohydrate),  has  the 
specific  gravity  1.8384  at  15°  (water  of  4°  = i),  and  is,  therefore,  lighter 
than  slightly  hydrous  acid.  When  the  anhydrous  acid  is  heated,  white 
fumes  of  sulphur  trioxide  escape  at  40°  ; the  liquid  begins  to  boil  at  200°, 
and  at  338°  the  acid,  with  1.5  per  cent,  of  water,  again  distils  over. 

From  these  data  it  is  obvious  that  sulphuric  acid,  even  at  a gentle  heat,  sustains  a par- 
tial decomposition  (dissociation)  into  sulphur  trioxide  and  -water.  The  vapor  density  of 
sulphuric  acid  has  been  found  to  be  72.0  (O.^  = 32)  at  332°  (near  its  boiling  point).  It 
diminishes  at  higher  temperatures,  and  is  49  at  416°,  where  it  is  constant.  This 
behavior  is  explained  by  the  dissociation  of  the  acid  molecules,  according  to  the  equation  : 

H2SO,  = SO,  + H2O. 

I vol.  I vol.  I vol. 

Hence  the  dissociation  of  the  acid  is  complete  at  416°,  while  it  is  only  about  34  per 
cent,  at  332°  (p.  94). 

Concentrated  sulphuric  acid  is  a thick,  oily  liquid.  On  cooling  a sul- 
phuric acid,  containing  about  15  per  cent,  of  water,  to  0°,  large  six-sided 
prisms  of  the  hydrate  H2SO^-|-H20  (called  sulphuric  acid  dihydrate) 
separate;  these  melt  at  H-8°,  and  at  205°  break  down  into  the  anhydrous 
acid  and  water.  The  second  hydrate,  H2SO4 2H2O,  has  the  specific 
gravity  1.63,  and  yields  water  at  195°.  The  concentrated  acid  possesses 
an  extremely  great  affinity  for  water,  and  absorbs  aqueous  vajior  energet- 
ically, hence  is  applied  in  the  drying  of  gases  and  in  desiccators.  It  mixes 
with  water  with  the  evolution  of  considerable  heat,  and,  for  this  reason,  it 
is  especially  directed,  in  diluting  the  acid,  to  pour  the  latter  in  a thin 
stream  into  the  water,  and  not  the  reverse,  as  explosive  phenomena  occur. 
In  mixing  suljihuric  acid  with  water,  a contraction  of  the  mixture  takes 


SULPHURIC  ACID. 


193 


place;  its  maximum  corresponds  to  the  hydrate  H2SO^-j-2H20  (/.  e., 
98  parts  : 36  parts). 

The  existence  of  the  hydrates  of  sulphuric  acid  is  explained,  as  in  the  case  of  periodic 
acid,  by  the  assumption  of  hydroxyl  groups.  Later  investigations  confirm  this  idea  : 

H2S0^  + = S(OH)g,  Hexahydroxyl  sulphuric  acid. 

11280^  -(-  11.2^  — SO(OH)^,  Tetrahydroxyl  sulphuric  acid. 

11280^  . . = S02(0H)2,  Normal  sulphuric  acid. 

The  tetrahydroxyl  and  the  hexahydroxyl  sulphuric  acids  yield  only  salts  of  the  normal 
dibasic  acid,  when  they  are  acted  upon  by  bases.  8alts  corresponding  to  the  hydrates  are 
not  known,  as  was  the  case  with  periodic  acid. 


The  affinity  of  sulphuric  acid  for  water  is  so  great  that  the  former  with- 
draws the  hydrogen  and  oxygen  from  many  substances,  with  the  produc- 
tion of  water.  In  addition  to  carbon,  many  organic  compounds  contain 
hydrogen  and  oxygen  in  the  proportion  in  which  these  elements  yield 
water.  The  withdrawal  of  hydrogen  and  oxygen  from  such  substances 
leaves  the  carbon.  This  explains  the  charring  action  of  the  acid  upon 
wood,  sugar,  and  paper.  When  sulphuric  acid  acts  upon  alcohol  (C2HgO), 
ethylene,  C2H^,  results  (p.  153). 

By  conducting  sulphuric  acid  vapors  over  red-hot  porous  bodies,  it  is 
decomposed  into  sulphur  dioxide,  water,  and  oxgen  (p.  194)  : 


H28O,  SOj  + lIjO  + O. 

When  heated  with  sulphur,  phosphorus,  carbon,  and  some  metals  (mer- 
cury, copper),  the  acid  is  reduced  to  sulphur  dioxide  (p.  183).  Nearly 
all  the  metals  are  dissolved  by  it,  forming  salts;  only  lead,  platinum,  and 
a few  others  are  scarcely  attacked  at  all.  It  is  a very  strong  acid,  and, 
when  heated,  expels  most  other  acids  from  their  salts;  upon  this  depends 
its  application  in  the  manufacture  of  hydrochloric, nitric,  and  many  other 
acids,  especially  the  organic  acids.  The  barium  salt  (BaSOJ  is  character- 
ized by  its  insolubility  in  water,  acids,  and  alkalies;  therefore,  sulphuric 
acid  added  to  solutions  of  barium  compounds  produces  a white  pulveru- 
lent precipitate,  which  serves  to  detect  small  quantities  of  the  acid. 

The  structure  of  sulphuric  acid  and  its  anhydride  can,  assuming  sul- 
phur to  be  sexivalent,  be  expressed  by  the  following  formulas: 


VI 

8 O. 


Pyrosulphuric  or  Disulphuric  Acid,  H2S20^.  On  withdrawing 
one  molecule  of  water  from  two  molecules  of  the  acid  there  results  the 


compound  1^2820^,  whose  formation  and  structure  may  be  represented  by 
the  following  formula : 


SO 

SO 


- H20 


S02< 

S02< 


OH 

O 

OH. 


As  this  acid  contains  two  hydroxyl  groui)s  it  is  dibasic;  yet  its  manner 
of  formation  shows  that  it  j)ossesses  an  anhydride  character.  Later,  we 

17 


94 


INORGANIC  CHKMISTRY. 


will  observe  that  almost  all  polybasic  acids,  like  i)hosphoric  acid, 
PC)(0H)3,  silicic  acid,  Si()(OII)2,  and  chromic  acid,  Cr()2(()n)2,  arc 
capable,  by  the  condensation  of  several  molecules  and  the  elimination  of 
water,  of  forming  like  derivatives,  which  bear  the  name  Poly-  or  Pyro- 
acids.  The  pyro-acid,  corresponding  to  sulphurous  acid,  is  also  known 
in  salts. 

The  disulphuric  acid  is  contained  in  the  so-called  fuming:;  or  Nordhau- 
sen  sulphuric  acid  (Acidum  sulphuricum  fumans,  oil  of  vitriol),  which  was 
formerly  obtained  by  heating  dehydrated  ferrous  sulphate — green  vitriol 
(FeSO^).  It  is  a thick,  oily,  strongly  fuming  li([uid,  of  specific  gravity 
1.86-1.9.  When  it  is  cooled,  large  colorless  crystals  of  disulphuric  acid 
(H2S2O2)  separate ; these  melt  at  35°.  Heat  breaks  it  down  into  sulphuric 
acid  and  suli)hur  trioxide,  which  volatilizes: 

II2S2O7  = up(\  -f  SO3. 

Conversely,  disulphuric  acid  may  be  obtained  by  dissolving  sulphur 
trioxide  in  sul})huric  acid  : 

II2SO,  + SO3  = II2S2O7. 

The  production  of  fuming  sulphuric  acid  also  depends  on  this,  as  it  may 
be  regarded  as  a solution  of  sulphur  trioxide  or  pyrosuli)huric  acid  in  an 
excess  of  sulphuric  acid. 

Technically,  fuming  sulphuric  acid  is  obtained  from  pyrites  (FeS2) — (at  present  only 
in  Bohemia).  The  decomposition  of  the  pyrites  in  the  air  yields  ferrous  sulphate  and 
ferric  sulphate,  which  can  be  dissolved  out  with  water.  The  solution  is  evaporated,  and 
the  residue  roasted  in  a reverberatory  furnace,  whereby  the  ferrous  salt  is  changed  to 
ferric  salt.  The  latter  is  then  distilled  from  earthen  retorts,  when  sulphuric  acid  and  sul- 
phur trioxide  pass  over  and  are  collected  in  the  receivers  : 

Fe,(SO.)3  = Fe^O,  + 3SO3. 

The  residue,  consisting  of  red  ferric  oxide,  finds  application  as  colcothar  {caput  7nor- 
tuum)  in  polishing  and  as  a paint. 

Solid,  crystalline  pyrosulphuric  acid  has  been  recently  introduced  into 
the  market  as  a substitute  for  the  fuming  liquid  sulphuric  acid.  It  is  made 
by  conducting  the  theoretical  amount  of  sulphur  trioxide  into  concen- 
trated sulphuric-acid  (see  above).  Sulphur  trioxide  is  prepared  by  two 
distinct  methods  at  present. 

In  Winkler’s  method  sulphuric  acid  of  66°  Baume  is  first  allowed  to  run  into  retorts 
raised  to  a red  heat.  The  mixed  gases,  sulphur  dioxide,  oxygen  and  water  (p.  193  b 
resulting  from  this  action,  are  freed  from  steam  by  passing  through  a coke  tower  through 
which  trickles  concentrated  sulphuric  acid.  The  dry  mixture  is  then  conducted  over 
heated  jdatinized  balls  of  white  clay  and  the  resulting  sulphur  trioxide  is  collected  in 
concentrated  suljfiniric  acid. 

'I'lie  more  recent  method  of  Wolters  consists  in  ]>roducing  sodium  pyrosulphate  by  heat- 
ing sodium  sulphate  with  concentrated  sulphuric  acid  : 


Na2SC),  -j-  ILSOi  -f  II./X 

Sodium  Sodiiiin 

sulphate.  pyrosulphate. 


SULPHURIC  ACID. 


95 


An  intermediate  product  in  this  reaction  is  primary  sodium  sulphate — NallSO^ — 
which  upon  the  application  of  heat  gives  up  water  and  passes  into  the  pyrosulphate.  By 
the  action  of  concentrated  sulphuric  acid  upon  the  latter  the  anhydride  is  liberated  and 
distilled  off  in  a vacuum  : 

Na^S.O,  + H^SO^  = 2NaHS04  + SO3. 

The  residual  sodium  sulphate  can  again  be  converted  into  pyrosulphate.  Sodium  pyro- 
sulphate decomposes  at  about  600°  into  sodium  sulphate  and  sulphur  trioxide,  thus  : 

Na2S207  = Na2S04  -j-  SO3. 


Sulphuric  Acid  Chloranhydrides. — Under  the  name  of  halogen  anhydrides  we 
understand  the  derivatives  resulting  from  the  replacement  of  hydroxyl  in  acids  by 
chlorine.  Conversely,  the  chloranhydrides,  by  the  action  of  water,  pass  into  the  corre- 
sponding acids  ; 

so,<^}  + 2H,0  = + 2HCI. 

The  ordinary  method  for  the  preparation  of  the  chloranhydrides  consists  in  permitting 
phosphorus  pentachloride  to  act  on  the  acids.  Sulphuric  acid  has  two  hydroxyl  groups  ; 
therefore  it  can  furnish  two  chloranhydrides. 

Cl 

The  first,  S02<^j^q,  Sulphuryl  Hydroxy-chloride,  or  Chlorstdphonic  Acid,  results 
when  one  molecule  of  phosphorus  pentachloride  acts  upon  one  molecule  of  sulphuric  acid  : 

+ PCI5  = SO,<g'j^  + FOCI,  + HCl. 

The  resulting  oxychloride  of  phosphorus  acts  upon  two  additional  molecules  of  sul- 
phuric acid,  with  formation  of  metaphosphoric,  hydrochloric,  and  chlorsulphonic  acids  : 

2S02(0H)2-f  POCI3  = 2S02<qj^  + HPO3  + HCl. 

It  is  formed,  too,  by  the  direct  union  of  sulphuric  anhydride  with  hydrochloric  acid : 
SO3  -f  HCl  = SO3CIH. 

The  mo.st  practical  method  for  its  formation  consists  in  conducting  chlorine  gas  through 
sulphuric  acid  (15  parts),  and  gradually  adding  phosphorus  trichloride  (7  parts).  Or, 
hydrochloric  acid  gas  is  led  into  solid  fuming  sulphuric  acid  (H2S2O.J),  as  long  as  absorp- 
tion occurs,  and  then  the  acid  is  distilled  (Otto). 

Chlorsulphonic  acid  is  a colorless,  strongly  fuming  liquid,  of  specific  gravity  1.716  at 

18°,  and  boils  without  decomposition  at  152°.  The  salt  S02-<qj^  results  from  the 

union  of  sulphur  trioxide  with  potassium  chloride. 

The  second  chloranhydride,  SO2CI2,  or  Sulphury!  Chloride,*  forms  when  phosphorus 
pentachloride  acts  upon  sulphuric  anhydride  ; and  also  by  the  direct  union  of  sulphur  di- 
oxide with  chlorine  in  sunlight : 

SO2  -(-  Clj  = SO2CI2. 

I vol.  I vol.  I vol. 


* 'Hie  group  SO2  combined  with  20H-groups  in  sulphuric  acid,  is  known  as  Sul- 
phuryl, 


96 


INORCJANIC  CHEMISTRY. 


'I’he  most  convenient  method  for  its  formation  is  to  pass  equal  volumes  of  sulphur  dioxide 
and  chlorine  gas  into  a cajxacious  flask,  containing  some  camphor,  which  causes  the 
union  of  the  gases  to  form  sulphuryl  chloride.  A colorless,  suffocating,  strongly  fuming 
liquid,  of  specific  gravity  1.667  at  20°  (referred  to  water  at  4°),  results.  It  boils  at 
69. 1°.  Water  decomposes  it  slowly  at  the  ordinary  temperature  into  sulphuric  and  hydro- 
chloric acids.  A little  water  will  first  change  it  to  chlorsulphonic  acid  : 

SOj<°[  + 11,0  = SO,<°'jj  + IICI. 


Its  vapor  density  is  normal  at  130°  ; at  higher  temperatures  it  breaks  down  into  sulphur 
dioxide  and  chlorine  : 


SO2CI2  = SO2  + CI3. 


Pyrosulphuryl  Chloride,  S.^O^Clj,  is  the  chloranhydride  of  pyrosulphuric  acid.  It 
is  obtained  by  several  reactions,  chiefly,  however,  by  the  action  of  phosjdiorus  pentachlo- 
ride  or  phosphorus  pentoxide  upon  chlorsulphonic  acid  : 


It  is  a thick  liquid,  fuming  in  the  air;  has  a specific  gravity  of  1.819  at  18°,  and 
boils  at  142°.  At  210°  it  shows  a normal  density  ; it  is  dissociated  at  higher  tempera- 
tures, corresponding  to  the  equation  : 


S2O5CI2  SO3  + SO2  + CI2. 

It  dissolves  gradually  in  water,  without  hissing,  and  decompo.ses  into  sulphuric  acid  and 
hydrochloric  acid  ; at  first,  with  a little  water  it  yields  chlorsulphonic  acid. 


AMIDO-DERIVATIVES  OF  SULPHURIC  ACID. 

The  amide  of  an  acid  is  produced  when  its  hydroxyl  is  replaced  by  the  univalent 
group  NH2.  In  the  case  of  polybasic  acids  there  are  several  amides  as  there  are  sev- 
eral chlorides,  depending  upon  whether  the  hydroxyl  groups  are  partly  or  completely  re- 
placed by  amido-groups.  Sulphuric  acid  yields  : 


S02< 


NH 

OH 


3 


Amidosulphonic  acid, 
Sulphanilic  acid 
or  Amidosulphuric  acid. 


SO 

^^2^NH2 

Sulphamide, 
Sulphuric  acid  amide. 


They  may  be  obtained  by  treating  the  corresponding  chlorides  with  dry  ammonia. 

Siilphamide,  formed  by  Traube  on  conducting  ammonia  into  a solution  of  sulphuric 
acid  chloride  : 

SO2CI2  + 4NH3  = S02(NH2)2  + 2NH,C1, 


consists  of  large,  colorle.ss  crystals  melting  at  8l°.  It  dissolves  readily  in  water  ; its  solu- 
tion is  neutral  in  reaction  and  has  a .slightly  bitter  taste.  Alkalies  and  acids  partly 
.saponify  it  in  acjueous  solution,  forming  at  first  sulphamic  acid,  then  sulphuric  acid  and 
a\nmonia — a reaction  corresponding  to  the  decomposition  of  the  chlorides  of  sulphuric 
acid  by  water.  Suli)hamide,  like  sulphuric  acid,  contains  two  hydrogen  atoms  replaceable 
by  metals.  It  yields  a silver  salt  of  the  composition  S02(NIIAg)2  [see  W.  Traube,  Ber. 
26  (1893),  I,  607]. 

In  the  action  of  ammonia  upon  sulphuryl  chloride,  sulphimide,  (SOg)  = Nil,  is  pro- 
d iced  together  with  suljdiamide.  It  is  also  formed  when  the  latter  is  heated  to  200-210°. 
It  acts  like  a monobasic  acid  ; its  salts,  e.  g.^ 

SO2N.K,  S02N.Na,  (S02N)2Ba -f  2H2O, 


POLYTHIONIC  ACIDS. 


197 


are  readily  soluble  in  water  and  crystallize  well.  Crystallized  sulphiinide  is  not  yet 
known  ; by  the  absorption  of  water  it  passes  readily  into  sulphuric  acid  and  ammonia 
(\V.  Traube). 

In  addition  to  the  preceding  there  are  other  compounds  derived  either  from  ammonia 
or  from  hydroxylamine  in  such  manner  that  the  hydrogen  in  them  in  direct  union  with 
nitrogen  is  replaced  by  the  univalent  sulpho-group — SO3H.  Two  series  result : 


NH^.  SO3H 

Amidosulphuric  acid  (see  p.  196). 

NHISO^H)^ 

Imidosulphonic  acid, 
Disulphamic  acid. 

N(S03H)3 

Nitrolosulphonic  acid, 
Trisulphamic  acid. 


HO.  NH.  SO3H 

Hydroxylamine  sulphonic  acid. 

HO.  N(S03H)2. 

Hydroxylamine  disulphonic  acid. 


The  acids  of  the  first  series  are  obtained  from  ammonia  and  sulphur  trioxide  or  chlorsul- 
phonic  acid  ; also,  just  like  the  second,  from  sulphurous  acid  and  nitrous  acid  or  hydroxyl- 
amine in  aqueous  solution  [Raschig,  Ann.  Chem.  241  (1887),  161  ; also  Chem.  Central- 
blatt,  1897,  I,  10 ; compare  p.  131]. 

OH 

Thiosulphuric  Acid,  H2S2O3  = ? generally  known  as  hypo- 

sulphurous  acid,  can  be  considered  as  sulphuric  acid  in  which  the  oxygen 
of  an  hydroxyl  group  is  replaced  by  sulphur.  It  is  not  known  in  a free 
condition,  since  as  soon  as  it  is  liberated  from  its  salts  by  stronger  acids 
it  at  once  decomposes  into  sulphur  dioxide,  sulphur  and  water  (com- 
pare Ber.  22  (1889),  1686,  2703): 

Na2S203  -f  2HCI  = 2NaCl  + SO2  -f  S -f  H2O. 

Its  salts,  called  hyposulphites,  are  of  practical  importance  (compare  Sodium 
Hyposulphite).  They  are  formed  by  the  direct  addition  of  sulphur  to 
sulphites  : 

Na2S03  S = Na2S203  ; 


similar  to  the  formation  of  sulphates  by  the  addition  of  oxygen  to  the  sul- 
phites. 


The  formation  of  its  sodium  salt  by  the  action  of  iodine  upon  a mixture  of  sodium  sul- 
phite and  sodium  sulphide, 

NaSNa  -f  NaS03 . Na  + I2  = NaS . SO.p^a  -f  2NaI, 

Sodium  hyposulphite. 

is  taken  as  evidence  of  the  chemical  structure  of  the  acid  (Spring).  However,  this  reac- 
tion can  take  place  in  two  phases  : 

I . Na2S  + B 2NaI  + S ; 2.  S + Na2S03  = Na2S203. 

Sodium  amalgam  resolves  sodium  hyposulphite  into  sodium  sulphite  and  sodium  sulphide. 


POLYTHIONIC  ACIDS. 

By  this  name  (from  tzoXuc;,  many,  and  Osiov,  sulphur)  are  understood  the 
acids  of  sulphur,  containing  two  or  more  atoms  of  the  latter  with  six 


198 


INORGANIC  CIIEMISIRY. 


oxygen  atoms  and  two  Iiydrogen  atoms.  Salts  of  the  following  are 
known : 

l^ithionic  acid. 
ll-^SgOg,  'rrithionic  acid. 
llgS^Og,  d'etrathionic  acid. 
ll.^SjOg,  rcMtathionic  acid. 


Following  Blomstrand  and  Mendelejeff  it  maybe  assumed  that  there  are  two  univalent 
\’I 

groups — SOgll  or  SOg.OH — in  the.se  acids  which  in  dilhionic  acid  are  directly  com- 
bined, but  in  the  other  acid.s  are  joined  through  S,  S.^,  S3,  with  one  another. 

SO3II  SO3II  SO3II  SO.II 

SO3II  'SO3II 

VI 

The  group  SO3II  is  known  as  the  sulpho-growY>  •,  it  is  also  present  in  organic  sulpho- 
acids,  and  corresponds  to  the  acid-forming  carbon  group,  COOII,  called  carboxyl.  From 
this  group  are  derived  the  acids  just  mentioned  : 


H.SOj.OH  HO.SOg.OII  ^<80  011  IIS.SOg.OH. 

Sulphurous  acid.  Sulphuric  acid.  Pyrosulphuric  acid.  Thiosulphuric  acid. 

The  formulas  of  the  polythionic  acids  may  also  be  derived  from  the  formulas  of  the 
hydrogen  sulphides,  HjS,  IlgSg,  and  HgSg,  if  the  hydrogen  atoms  in  the.se  be  replaced 
by  the  sulpho-group.  If  in  the  case  of  the  ordinary  hydrogen  sulphide  this  replacement 
extends  merely  to  one  hydrogen  atom,  thiosulphuric  acid  results  ; the  acids  corresponding 
to  it  are  not  known  in  the  other  two  series : 


c/H 

e .SO3  . OH 

CJ  .SOg.  OH 
^^SOg.  OH 

s<r^ 

• • 

C.  ^SO., . OH 
. OH 

Q ^SOg  . OH 
^3<-S0.,  .OIL 

Similarly,  sulphuric  acid  and  pyrosulphuric  acid  are  derived  from  water ; whereas 
persulphuric  acid  must,  according  to  Michaelis  and  Richarz,  be  referred  to  hydrogen 
peroxide  : 


H 

H 

Q^SOgH 

O.SO3H 

^^SOgH 

Sulphuric  acid. 

Pyrosulphuric  acid. 

-H 

H 

0 <-S03H 

Q ^SOgH 

Unknown. 

Persulphuric  acid. 

This  is  an  example  of  the  type  idea  mentioned  on  p.  1 72.  Whether  the  formulas 
deduced  above  actually  repre.sent  the  true  chemical  structure  of  the  polythionic  acids  is 
at  present  very  questionable  [see  Debus,  Ann.  Chem.  244  (1888),  181]. 

Dithionic  Acid,  Il^S./Ig,  is  only  known  in  aqueous  solution.  When  concentrated 
in  vacuo  or  when  heated  it  decompo.ses  into  sulphuric  acid  and  sulphur  dioxide.  Its 
manganc.se  .salt  (together  with  manganous  sulphate)  results  from  the  action  of  sulphur 
dioxide  upon  manganese  dioxide  (MnOj)  suspended  in  water: 

MnO.,  + 2SO.,  = MnS-Pg 

and 

MnOg  }-  SOg  = MnSO^. 


OXYGEN  DERIVATIVES  OF  SELENIUM  AND  TELLURIUM.  I99 

Barium  hydroxide  converts  this  into  the  barium  salt,  from  which  the  free  dithionic  acid  in 
aqueous  solution  is  obtained  by  means  of  sulphuric  acid. 

Trithionic  Acid,  HgSgOg,  is  not  known  in  a free  condition.  Its  salts  are  produced 
when  an  aqueous  solution  of  primary  potassium  sulphite  is  digested  with  flowers  of 
sulphur  : 

6KHSO3  -f  2S  = 2K2S3O6  + K2S2O3  + 3H,0. 

Separated  from  its  salts  by  other  acids  it  decomposes  rapidly  into  sulphuric  acid,  sulphur 
dioxide,  and  sulphur.  Its  production  by  the  action  of  iodine  upon  a mixture  of  sodium 
sulphite  and  hyposulphite  is  especially  interesting  : 

Na^SjOg  -j-  Na^SOj  -|-  I^  = NagSjOg  -f  2NaI. 

It  is  also  formed  from  sulphur  dichloride  and  sodium  sulphite  : 

SCI3  + 2Na2S03  — Na^SgOg  -}-  2NaCl  (Spring). 

The  tri-  and  not  a tetrathionate  is  produced  when  sulphur  dichloride  is  used. 

Tetrathionic  Acid,  H^S^Og.  Its  salts  are  produced  when  iodine  acts  upon  solutions 
of  the  hyposulphites  : 

2KS  . SO3K  4-  I2  = K,S,Og  + 2KI. 

Potassium  tetrathionate. 

This  reaction  is  applied  in  volumetric  analysis  for  the  quantitative  determination  of 
iodine,  and  such  substances  as  separate  iodine  from  potassium  iodide  (see  Sodium  Thio- 
sulphate). 

'fhe  tetrathionic  acid,  separated  from  its  salts  by  stronger  acids,  is  very  unstable,  and 
when  its  aqueous  solution  is  concentrated  it  decomposes  into  sulphuric  acid,  sulphur 
dioxide,  and  sulphur.  An  aqueous  solution  of  the  acid  can  be  easily  prepared  by  con- 
ducting hydrogen  sulphide  into  aqueous  sulphurous  acid  : 

4SO2  -|-  3^2^  — 2^2^  “f"  3^- 

If  the  solution  be  saturated  with  bases,  neutral  and  acid  tetrathionates  result,  e,  g.y 
BaS^Og  -f  2H2O  and  Zn(HS^Og)2. 

Pentathionic  Acid,  H2S50g,  is  supposed  to  be  produced  together  with  tetrathionic 
acid  when  hydrogen  sulphide  acts  upon  aqueous  sulphurous  acid  : 

5H2S  + 5SO2  - H2SgOg  + 4H2O  + 5S 

(Wackenroder’s  liquid  ; the  dry  gases  do  not  react).  It  was  from  this  solution  that  Debus 
first  prepared  pure,  crystallized  pentathionates. 

All  the  polythionic  acids  are  distinguished  from  sulphuric  acid  by  the  solubility  of  their 
barium  salts.  The  alkali  salts  of  pentathionic  acid  decompose  in  aqueous  solution  in 
such  a manner  that  after  some  time  the  salts  of  the  other  polythionic  acids  are  also  present 
in  the  solution. 


2.  OXYGEN  DERIVATIVES  OF  SELENIUM  AND  TELLURIUM. 

Se02  H2Se03 

Selenium  dioxide.  Selenious  acid. 

. . . IlgSeO,. 

Selenic  acid. 

Selenium  Dioxide,  Se02,  or  selenious  anhydride,  is  produced  when 
selenium  burns  in  the  air  or  in  oxygen.  It  consists  of  long  white  needles, 
which  sublime  at  about  320°  without  fusing.  It  dissolves  readily  in 


200 


INORGANIC  CHEMISTRY. 


water,  forming  selenious  acid,  H^SeO,.  Tlie  latter  is  also  obtained  by 
dissolving  the  metal  in  concentrated  nitric  acid.  When  the  solution  is 
evaporated  it  crystallizes  in  large,  colorless  jirisms,  which  decompose,  on 
heating,  into  the  anhydride  and  water.  Sulphurous  oxide  reduces  seleni- 
ous acid,  with  separation  of  free  selenium  : 

HjSeOj  -f  2SO2  + 11,0  = 2lI,SO,  -f  Se. 

It  is  similarly  reduced  by  potassium  iodide  in  hydrochloric  acid  solu- 
tion with  the  simultaneous  sejiaration  of  iodine: 

SeO,  + 4KI  + 4IICI  = 4KCI  + 2ll,0  4-  Se  -f  2I,. 

Selenic  Acid,  H2SeO^,  is  obtained  by  conducting  chlorine  gas  into 
an  acpieous  solution  of  selenious  acid  : 

H2Se03  -f  II2O  + CI2  = IljSeO^  + 2IICI. 

The  solution  may  be  concentrated  until  it  attains  a specific  gravity  of  2.6  when  it 
becomes  an  oily  liquid,  similar  to  suljdiuric  acid,  and  contains  95  per  cent,  of  selenic  acid, 
n2Se04.  If  the  solution  be  heated  above  260°,  the  acid  breaks  down  into  selenium 
dioxide,  oxygen  and  water.  It  was  from  this  solution  that  (.larneron  and  Macallan 
obtained  the  pure  acid  as  a crystalline  mass,  melting  at  58°.  It  forms  a hydrate  with 
water,  H2SeO^ . H2O,  melting  at  25°.  Its  anhydride  is  not  known. 

The  salts  of  selenic  acid  are  known  as  selenates,  those  of  selenious  acid  as  selenites 
[see  Jahr.  Chem,  1889,  388]. 


The  derivatives  of  tellurium  are  very  similar  to  those  of  selenium. 

Tellurium  Dioxide,  TeO,,  results  when  tellurium  is  burned,  and  forms  a white 
crystalline  mass,  fusing  at  a red  heat  and  subliming.  It  is  almost  insoluble  in  water. 

Tellurous  Acid,  H2Te03,  is  produced  when  the  metal  is  dissolved  in  concentrated 
nitric  acid.  'Water  will  precipitate  it  from  such  a solution  in  the  form  of  a white  amor- 
phous powder.  On  warming,  it  readily  decomposes  into  the  dioxide  and  water. 

Telluric  Acid,  H.2TeO^.  Potassium  tellurate  is  produced  when  tellurium  or  its  di- 
oxide is  fused  with  saltpeter.  Tellurium  is  added  to  an  excess  of  dilute  nitric  acid  in 
which  it  dissolves  with  violent  reaction.  The  resulting  dioxide  is  oxidized  by  means  of 
a little  more  than  the  calculated  quantity  of  chromic  acid : 

STeO,  + 2Cr03  +3H2O  = 3H2Te04  -f  Cr203. 

The  greater  portion  of  the  telluric  acid  separates  on  evaporation.  It  is  purified  by  pre- 
cipitation with  concentrated  nitric  acid  from  its  aqueous  solution  [Z.  f.  anorg.  Ch.  X 
(1895),  190],  The  acid  crystallizes  from  its  aqueous  solution  with  two  molecules  of 
water  (Il2Te04  -f-  211,0)  in  isometric  or  triclinic  forms,  which  at  100°  lose  all  their 
water  and  break  down  into  a white  powder  of  telluric  acid,  H2Te04.  The  latter  is  not 
very  soluble  in  water,  and  manifests  a slight  acid  reaction.  When  carefully  heated, 
telluric  acid  breaks  down  into  water,  and  the  trioxide,  TeOg,  which  is  a yellow  mass  in- 
soluble in  water,  and  by  further  application  of  heat  decomposes  into  the  dioxide  and 
oxygen. 


3.  OXYGEN  DERIVATIVES  OF  THE  ELEMENTS  OF  THE 
NITROGEN  GROUP. 

'Phc  halogens  combine  with  one  atom  of  hydrogen  to  form  halogen 
hydrides;  their  oxygen  acids  also  containing  one  atom  of  hydrogen. 


OXYGEN  DERIVATIVES  OF  NITROGEN. 


201 


The  elements  of  the  sulphur  group  contain  two  atoms  of  hydrogen  in 
their  hydrogen  derivatives  and  oxygen  acids.  Accordingly,  we  find  that 
the  elements  of  the  nitrogen  group  combine  with  three  atoms  of  hydro- 
gen, and  form  acids  which  also  contain  three  atoms  of  the  same  element: 

HCl  H^S  PHg 

HCIO3  H2SO3  P0,H3 

HCIO,  H^SO^  P0,H3. 


The  acids  of  the  nitrogen  group  containing  three  atoms  of  hydrogen, 
designated  normal  or  ortho-acids  {opdoq,  correct,  real,  true),  as  HgPO^, 
H3AsO^,  H3ASO3,  can  yield  monobasic  acids  by  the  removal  of  one  mole- 
cule of  water.  Such  derivatives,  having  one  atom  of  hydrogen,  are  called 
jneta-acids  (^era,  a word  which  here  expresses  a change  in  condition)  : 


H3PO, 

Orthophosphoric  acid, 
Na3As03 

Sodium  orthoarseniate. 


HPO3 

Metaphosphoric  acid. 

NaAsO.^. 

Sodium  meta-arseniate. 


These  meta-acids  of  phosphorus  pass  into  the  ortho-acids  by  the  absorp- 
tion of  water.  The  ortho-acids  of  nitrogen,  on  the  other  hand,  are  less 
stable  and  only  exist  in  some  salts.  The  ordinary  acids  and  salts  of 
nitrogen  belong  to  the  meta-series  and  contain  one  atom  of  hydrogen 
(or  metal) ; 

(H3NOJ  HNO3 

Orthonitric  acid.  Ord.  Nitric  acid. 


(H3NO3) 
Orthonitrous  acid. 


HNO2. 

Ord.  Nitrous  acid. 


The  further  exit  of  water  produces  the  true  anhydrides  (p.  183). 


1. 


OXYGEN  DERIVATIVES  OF  NITROGEN. 


N205  



HNO3 

Nitrogen  pentoxide. 

N.2O4  and  NOj  — 
Nitrogen  Dioxide, 

tetroxide. 

Nitric  acid. 

XT- 

JNitrogen  tnoxide. 

NO 

Nitric  oxide. 


> HNO2 
Nitrous  acid. 


N,0  ^ 

Nitrous  oxide. 


> (HNO).,. 

Hyponitrous  acid. 


The  following  formulas  express  the  structure  of  these  compounds: 


III  III 
N=N 
V 
o 

Nitrous  oxide. 


Ill  III 

0-=N  N^O 

V 

o 

Nitrogen 

trioxide. 


Ill  III  III 

TIO-N=N-OTI  OrrN-OII 

Hyponitrous  acid.  Nitrous  acid. 


Ill  V 
0=N 

V 

o 


Nitric-nitrous 

anhydride 

(tetroxide). 


V V 

V 

o 

Nitrogen 

pentoxide. 


O-N-OU. 
Nitric  acid. 


202 


INORGANIC  CHEMISTRY. 


Tlie  salts  of  nitric  acid  arc  called  ^titrates ; those  of  nitrous  acid, 
nitrites. 


NITRIC  ACID. 

IINO3. 

This  acid  occurs  in  nature  only  in  the  form  of  salts, — potassium, 
sodium,  and  calcium  nitrates  (comj)are  these), — which  have  resulted  from 
the  decay  of  nitrogenous  organic  substances  in  the  j^resence  of  strong 
bases  (the  alkalies)  and  Bacillus  nitrificans.  It  is  sometimes  present  in 
the  air  as  ammonium  salt.  The  free  acid  is  formed  in  very  slight  quan- 
tity by  conducting  electric  sparks  through  moist  air;  most  readily  with 
a mixture  of  5 volumes  of  air  and  6 volumes  of  oxygen,  which  may 
obtain  technical  value  if  the  electricity  and  oxygen  can  be  prepared 
cheaply  enough.  At  present  Chile  saltpeter  is  the  j)rincii)al  source  of 
nitric  acid. 

To  prepare  nitric  acid  heat  potassium  or  sodium  nitrate  in  a retort  with 
sulphuric  acid,  when  the  nitric  acid  will  distil  over  and  acid  sodium  or 
potassium  sulphate  remain  : 

NaNOa  -f  HjSO^  = IINaSO^  + IINO3. 

If  one  molecule  of  sulphuric  acid  be  used  with  two  molecules  of  the 
nitrate,  the  resulting  acid  sulphate  will  at  higher  temperatures  act  upon 
the  second  molecule  of  nitrate  in  the  sense  of  the  equation  : 

NallSO,  + NaNOa  = HNO3  -f  Na^SO,. 

The  temperature  requisite  to  complete  the  reaction  is,  however,  so  high 
that  a portion  of  the  nitric  acid  will  be  decomposed. 

Perfectly  anhydrous  nitric  acid  has  not  yet  been  obtained.  The  most 
concentrated  acid  (99.8  per  cent.  HNO3)  is  a colorless  liquid  of  specific 
gravity  1.56  at  0°  ; it  fumes  in  the  air,  and  at  — 47°  solidifies  to  a crys- 
talline mass.  At  ordinary  temperatures  it  undergoes  a partial  decomposi- 
tion (similar  to  sulphuric  acid)  into  water,  oxygen,  and  nitrogen  dioxide, 
NO2,  which  dissolves  in  the  acid,  with  a yellowish-brown  color;  the  color- 
less acid  therefore  becomes  colored  upon  standing,  and  in  sunlight  soon 
turns  yellow.  At  86°  the  acid  commences  boiling  and  sustains  a partial 
decomposition  ; the  first  portions  are  colored  yellow  by  the  dissolved 
nitrogen  dioxide,  but  subsequently,  aqueous  acid  distils  over.  Nitric 
acid  is  completely  decomposed  into  nitrogen  dioxide,  oxygen,  and  water, 
when  its  vapors  are  heated  to  about 

2IINO3  = 2NO2  + H3O  -f  O. 

The  acid  mixes  in  all  proportions  with  water.  Upon  distilling  the 
dilute  aqueous  solution,  only  pure  water  passes  over  at  first ; the  boiling 
temperature  gradually  rises,  and  between  120°  and  121°  a solution  goes 
over,  which  contains  68  ])er  cent,  of  HNO3,  and  has  a s])ecific  gravity  of 
1. 414  at  15°.  Tliis  is  the  ordinary  concentrated  nitric  acid  of  trade. 
When  this  is  distilled  with  5 parts  of  sul})huric  acid,  an  almost  anhydrous 
acid  is  obtained,  which  may  be  freed  from  nitrogen  dioxide,  contained  in 
it,  by  conducting  a stream  of  air  through  it. 


NITRIC  ACID.  . 203 

The  liquid  boiling  at  121°,  however,  can  be  regarded  as  a mixture  of  the  trihydrate, 
N0(0H)3,  and  pentahydrate,  N(OH)5  (pp.  193,  201). 

Nitric  acid  is  a very  powerful  acid,  oxidizing  or  dissolving  almost  all 
metals  (gold  and  platinum  excepted).  Nearly  all  the  metalloids,  like 
iodine  sulphur,  phosphorus,  and  carbon,  are  converted  by  it  into  their 
corresponding  acids.  It  acts  as  a very  strong  oxidizing  agent,  destroy- 
ing organic  coloring  substances,  and  decolorizes  a solution  of  indigo  very 
readily.  In  so  doing  the  nitric  acid  itself  is  deoxidized  to  the  lower 
oxidation  products  of  nitrogen  (NO  and  NOj).  Some  substances  even 
reduce  the  acid  to  ammonia.  Thus,  for  example,  if  zinc  be  brought  into 
dilute  nitric  acid  (5-6  per  cent.)  the  metal  will  be  dissolved  without  the 
liberation  of  hydrogen.  The  latter,  in  stain  nascendiy  acts  at  once  upon 
the  excess  of  acid  and  reduces  it  to  ammonia,  which  forms  an  ammonium 
salt  with  the  acid ; hence,  in  solution,  we  have  ammonium  nitrate  in 
addition  to  the  zinc  nitrate  : 

2HNO3  + Zn  = Zn(N03)2  + H, 

and 

2HNO3  + 4H2=  NH^N03  + 3H2O. 

If  the  aqueous  nitric  acid  be  less  dilute  (containing  more  than  10  per 
cent,  of  HNO3)  it  will  be  reduced  by  zinc  and  other  metals,  not  to  am- 
monia, but  to  the  nitrogen  oxides,  N2O,  NjOg,  and  N20^.  The  more 
concentrated  the  acid,  the  higher  will  the  oxides  be. 

The  reduction  of  nitric  acid  to  ammonia  by  nascent  hydrogen  occurs 
more  easily  in  alkaline  solution.  If  a solution  of  nitrates,  made  alkaline 
with  sodium  or  potassium  hydroxide,  be  heated  with  zinc  or  aluminium 
filings  or  iron  powder,  then  all  the  nitrogen  of  the  nitric  acid  will  be 
converted  into  ammonia : 

HNO3  + 4H2  = NH3  + 3H2O. 

Hydroxylamine  (p.  130)  and  ammonia  are  produced  when  nitric  acid 
acts  on  tin. 

Nitric  acid  usually  forms  salts  of  the  form  MeNOg,  with  one  equivalent 
of  the  metals;  these  are  called  nitrates,  and  are  all  readily  soluble  in  water. 

Red  Fuming  Nitric  Acid  {Acidum  nitricuni  fumans')  is  the  name  given 
a nitric  acid  containing  much  nitrogen  dioxide  in  solution.  It  is 
obtained  by  the  distillation  of  two  molecules  of  saltpeter  with  one  mole- 
cule of  sulphuric  acid  (p.  202),  or  better,  by  the  distillation  of  com- 
mercial concentrated  nitric  acid  with  concentrated  sulphuric  acid.  It 
generally  has  the  specific  gravity  of  i. 5-1. 54,  and  possesses  greater 
oxidizing  power  than  the  colorless  nitric  acid. 

A mixture  of  i volume  of  nitric  acid  and  3 volumes  of  concentrated 
hydrochloric  acid  is  known  as  Aqua  regia,  as  it  is  able  to  dissolve  gold 
and  platinum,  which  neither  of  the  acids  alone  is  capable  of  doing.  The 
powerful  oxidizing  action  of  ihe  mixture  is  due  to  the  presence  of  free 
chlorine  and  the  chlorine  derivative  NOCl,  which  may  be  considered 
the  chloranhydride  of  nitrous  acid  : 

3IICI  + IINO3  = 2II2O  + NOCl  -f-  CI3. 


204 


INORGANIC  CHEMISTRY. 


Nitrogen  Pentoxide,  N.^(  >5,  nitric  anliydi  ide,  is  produced  hy  carefully  heating  phos- 
phoric anhydride  with  nitric  acid  : 

211NO3 -f  I’A  - N/), -I  2iiro3; 

further,  together  with  oxygen,  on  conducting  chlorine  over  silver  nitrate  : 

2Ago.  NO,  + 2CI  = + o. 

It  forms  colorless,  rhombic  prisms,  melting  at  30°  and  boiling  with  partial  decomposi- 
tion at  47°.  It  is  very  unstable,  decomposing  readily  into  nitrogen  tetroxide  and  oxygen, 
and  sometimes  exploding  spontaneously.  It  yields  nitric  acid  with  water  and  evolves 
much  heat  by  the  union  : 

j;5g>>o  + ii,o  = 2NO,.on. 


Nitroxyl  Chloride,  NO.2CI,  the  chloranhydride  of  nitric  acid,  results  from  the  union 
of  nitrogen  dioxide  with  chlorine,  and  according  to  the  ordinary  method  of  forming  chlor- 
anhydrides  (see  p.  195),  by  the  action  of  phosphorus  pentachloride  or  oxychloride  upon 
nitric  acid,  or  better,  its  silver  salt : 

3NO, . OAg  + POCI3  = 3NO3CI  + P0(0Ag)3. 

Silver  nitrate.  Silver  phosphate. 

It  is  said  to  be  a heavy,  yellow  liquid,  boiling  at  -1-5°.  Water  immediately  decomposes 
it  into  hydrochloric  and  nitric  acids.  However,  Geuther’s  experiments  make  the  exist- 
ence of  this  acid  rather  doubtful  [Ann.  Chem.  245  (1888),  98]. 

Nitrosyl  Chloride,  NOCl,  is  produced  by  the  union  of  nitric  oxide  (2  volumes)  with 
chlorine  (l  volume):  2NO  CI2  = 2NOCI ; when  phosphoric  chloride,  PCI5,  is 
allowed  to  act  upon  liquid  nitrogen  tetroxide,  N2O4,  and  by  heating  lead-chamber  crystals 
(p.  208)  wdth  sodium  chloride  to  80-90°  : 

+ NaCl  = + NOCl. 


The  reddish -yellow  vapors  that  escape,  if  cooled  to  — 20°,  condense  to  a red  liquid  of 
specific  gravity  1. 416  at  — 12°,  and  boils  at  -|-2°.  It  forms  nitrous  and  hydrochloric 
acids  with  water  : 


NOCl  -f  H2O  = HNO2  + HCl. 


It  may,  therefore  be  regarded  as  the  chloranhydride  of  nitrous  acid — NO  . OH. 

O.  O. 

Nitranaide,  NO2 . NH2  = '^N  - NHj  or  ^N  = NH,  the  amide  of  nitric  acid, 

O^  HO^ 


was  di.scovered  by  Thiele  and  Lachmannon  decomposing  pota.ssium  nitrocarbaminate  (see 
Organic  Cliemistry)  with  sulphuric  acid.  It  is  also  produced  upon  adding  nitric  acid  to  a 
.solution  of  potassium  imidosulphonate  (p.  196)  in  concentrated  sulphuric  acid  : 


NH(S03H)2  + NO2.  OH  NO2.  NH2  + H2S2O,. 

Nitramide  crystallizes  from  ligroine,  in  which  it  dissolves  with  difficulty,  in  brilliant  white 
plates.  1 1 dissolves  readily  in  ether,  alcohol  and  water.  It  melts  at  about  73°  'vith  imme- 
diate decomposition  into  nitrous  oxide  and  water : 

NO2.  NH2  = N2O  4 IP/). 

It  volatilizes  partly  at  the  ordinary  temperature.  It  is  especially  sensitive  to  alkalies;  in 


NITROGEN  TRIOXIDE. 


205 


contact  with  sodium  hydroxide  it  explodes  with  fire  phenomena.  Its  aqueous  solution 
reacts  strongly  acid.  It  also  yields  a readily  decomposable  mercury  salt : 

NO^NHg. 

When  nitramide  is  reduced  hydrazine  results  : 

NO2NH2  + 3H2  = NHj . NHj  + 2H,0. 

Nitramide  is  isomeric  with  hyponitrous  acid  : HO-N=N-OH.  [See  Ann.  Chem.  288 
(1895),  267;  296  (1897),  95. J 


Nitrogen  Trioxide,  NjOg,  nitrous  anhydride,  is  formed  by  the 
direct  union  of  nitrogic  oxide  (4  volumes)  with  oxygen  (i  volume)  at 
—18°: 

4NO  + 02  = 2N2O3  ; 

4 vols.  I vol.  2 vols. 

by  mixing  liquid  nitrogen  tetroxide,  N20^,  with  a little  cold  water  : 

2 JJgqo  +hp  = noj.o  + 2N0,.0H; 

by  the  introduction  of  nitric  oxide  into  liquid  nitrogen  tetroxide 
below  — 21°  : 

N2O4  + 2NO  = 2N2O3 ; 

and  by  conducting  nitric  oxide  into  cold  anhydrous  nitric  acid : 

2HNO3  + 4NO  = 3N2O3  + HgO. 

It  is  most  easily  obtained  by  the  action  of  nitric  acid  upon  arsenious 
oxide,  AS2O3.  Nitrogen  tetroxide  is  simultaneously  produced,  from  which 
it  is  readily  separated  by  fractional  distillation  and  condensation. 

Nitrogen  trioxide  condenses  at  — 21°  to  a dark-blue  liquid  of  specific 
gravity  1.449  begins  to  boil  at  3.5°.  It  decomposes  when 

distilled ; its  vapors  consist  of  a mixture  of  the  tetroxide  and  nitric 
oxide  (N2O4  2NO)  and  at  more  elevated  temperatures  of  nitrogen  di- 
oxide and  nitric  oxide  (2NO2  -f-  2NO).  Upon  cooling  they  reunite  to 
the  liquid  nitrous  anhydride.  The  latter  is  therefore  only  known  in  the 
liquid  condition.  Its  dissociation  begins  at  — 21°  [Geuther,  Ann. 
Chem.  245  (1888),  96;  see  also  Lunge  and  Porschnew,  Z.  f.  anorg. 
Chem.  7 (1894),  209]. 

The  trioxide  mixed  with  a little  cold  water  probably  forms  nitrous 
acid  (N2O3  -f-  H2O  = 2HNO2) ; more  water,  aided  by  heat,  decomposes 
it  into  nitric  acid  and  nitric  oxide  gas : 

3HNO2  = HNO3  + 2NO  + H2O. 

Nitrous  Acid,  HNO2,  is  not  known  in  a free  state.  Its  salts  (the 
7iitrites)  are  obtained  by  igniting  the  nitrates: 

KNO^  = KNOj  + O. 


2o6 


INORGANIC  CHEMISTRY. 


The  withdrawal  of  oxygen  is  rendered  easier  if  oxidizalile  metals, 
lead,  be  added  to  the  fusion  (see  Potassium  and  Sodium  Nitrite).  Traces 
of  nitrites  are  found  in  the  air  and  in  many  waters. 

On  adding  sulphuric  acid  to  the  nitrites,  brown  vapors  are  disengaged  ; 
these  consist  of  nitrogen  dioxide  and  nitric  oxide.  It  may  be  that  the 
nitrous  acid,  at  first  liberated,  is  broken  down  into  water  and  the  trioxide, 
which,  as  we  have  seen  above,  gradually  decomposes  into  nitrogen  dioxide, 
NO2,  and  nitric  oxide,  NO.  Similar  reddish-brown  vapors  are  obtained 
if  nitric  acid  be  permitted  to  act  upon  starch  or  arsenious  oxide  (As.^O.,). 

According  to  lAinge,  if  we  einjdoy  nitric  acid  of  specific  gravity  1.30- 1.35,  nitrogen 
trioxide  is  produced  almost  exclusively,  whereas  in  using  the  concentrated  acid  (1.4-1. 5) 
we  get  a mixture  rich  in  dioxide,  and  if  the  acid  be  dilute  the  chief  product  is  nitric 
oxide,  NO,  with  a little  nitrogen  dioxide,  NOj. 

The  nitrous  acid  which  has  separated  out  in  the  solution  and  its  decom- 
position products — N()2  and  NO — are  strong  oxidizers,  setting  iodine 
free  from  the  soluble  iodides.  In  other  cases,  however,  they  exhibit  a 
reducing  action  ; thus,  e.  g.,  the  acidified  red  solution  of  jiotassium  jier- 
manganate  is  decolorized  by  the  addition  of  nitrites,  and  nitrous  becomes 
nitric  acid.  In  very  dilute  aqueous  solution,  the  action  proceeds  ac- 
cording to  the  following  equation  : 

5HNO2  + 2KMnO^  + 311280^  ==  5HNO3  4-  K2S0^  + 2MnS(\  3H2O. 

This  reaction  serves  for  the  quantitative  determination  of  free  nitrous 
acid,  as  well  as  for  its  salts  (p.  208). 


Nitrogen  Tetroxide,  N20^,  or  nitrogen  dioxide^  NO2  (formerly  called 
hyponitric  acid'),  only  exists  at  low  temperatures;  when  heated,  it  suffers 
a gradual  decomposition  into  the  simpler  molecules  NO2  which  recom- 
bine to  N20^  upon  cooling.  We  here  meet  the  interesting  case  of  disso- 
ciation, occurring  even  at  the  ordinary  temperature.  The  tetroxide, 
N20^,  is  colorless,  while  the  dioxide,  NO2,  is  colored  red-brown  ; it  appears, 
therefore,  that  the  color  gradually  becomes  darker  as  the  temperature 
rises,  and  that  it  corresponds  to  the  increasing  dissociation  of  the  com- 
plex molecules  N20^.  The  same  was  observed  with  sulphur  tetrachloride 
(p.  no)  and  nitrogen  trioxide  (p.  205). 

At  ordinary  temperatures,  nitrogen  tetroxide  is  a liquid  of  specific 
gravity  1.49.  When  cooled  to  — 20°  it  solidifies  to  a colorless  crystal- 
line mass,  melting  at — 12°.  In  consequence  of  the  dissociation  which 
begins  at  0°,  the  liquid,  at  first  colorless,  becomes  yellow,  and  the  in- 
tensity in  color  grows  with  rising  tem])erature.  The  liquid  begins  to 
boil  at  about  22°,  and  is  converted  into  a yellowish-brown  vapor  which 
becomes  dark  as  the  temiierature  is  increased. 

'flic  theoretical  vapor  density  of  nitrogen  tetroxide,  N20^  ecjuals  92.08,  while  that  of 
the  dioxide,  N()2,  ecpials  46.04.  'i'lie  exj)erimental  vaj)or  density  has  been  found  to  equal 
76  at  the  point  at  which  the  licpiid  compound  boils  (26°)  ; it  may  l)e  calculated  from  this 
that,  at  this  t(nn])cniture,  34.4  per  cent,  of  the  tetroxide  molecules  are  decomposed  into 
dioxide  molecules.  IJeiice  we  conclude  that  the  dissociation  of  the  compound  com- 


NITROGEN  TETROXIDE. 


207 


mences  even  in  the  liquid  state  ; this  is  confirmed  by  the  yellow  coloration  appearing  at  0°. 
With  rising  temperature  the  density  of  the  vapor  steadily  diminishes,  becomes  constant 
finally  at  150°  and  equals  46.  Then  all  the  molecules  (N2O4)  are  decomposed  into  the 
simpler  molecules  NO2  ; and  the  dark  coloration  of  the  vapors  attains  its  maximum. 

Nitrogen  tetroxide  is  produced  by  conducting  electric  sparks  through 
a mixture  of  dry  nitrogen  and  oxygen ; it  is  also  formed  by  the  union  of 
two  volumes  of  nitric  oxide  with  one  volume  of  oxygen  : 

2NO  + O2  = N2O,. 

2 vols.  I vol.  I vol. 

We  can  get  it  more  conveniently  by  heating  dry  lead  nitrate,  which 
decomposes  according  to  the  following  equation  : 

Pb(N03)2  = PbO  + O + 2NO2. 

The  escaping  dioxide  condenses  in  the  cooled  receiver  to  liquid 
tetroxide,  N20^. 

Nitrogen  tetroxide  varies  in  its  behavior  with  water,  according  to  the 
temperature  of  the  latter.  We  saw  that  by  the  action  of  a little  cold 
water,  it  was  decomposed  into  nitrogen  trioxide  and  nitric  acid  (p.  205). 
With  an  excess  of  cold  water,  and  also  with  an  aqueous  solution  of  alkalies, 
it  yields  nitric  and  nitrous  acids,  that  is,  their  salts: 

+ ^2^  = NO2 . OH  -L  NO  . OH. 

Both  reactions  plainly  indicate  that  the  liquid  tetroxide  represents  the 
mixed  anhydride  of  nitric  and  nitrous  acids  (p.  201)  ; similarly,  the  com- 
pound Cl20^  constitutes  the  mixed  anhydride  of  chloric  and  chlorous 
acids  (p.  1 1 7).  AVarm  water  converts  the  tetroxide  into  nitric  acid  and 
nitric  oxide — because  under  these  conditions  the  nitrous  acid  decom- 
poses into  nitric  acid  and  nitric  oxide  (see  p.  205).  The  dioxide  behaves 
similarly : 

3NO2  -f  H2O  = 2HNO3  + NO. 

The  tetroxide  and  dioxide  possess  strong  oxidizing  properties;  many 
substances  burn  in  their  vapors;  iodine  is  set  free  from  the  soluble 
metallic  iodides  by  them. 

Many  metals,  just  after  the  reduction  of  their  oxides  in  hydrogen,  absorb  large  quanti- 
ties of  nitrogen  dioxide  at  lower  temperatures,  e.  g.,  copper  takes  up  1000  times  its 
volume.  “ Nitro-metals  ” are  produced  in  this  way.  “ Nitro-copper,”  CiqNOg,  is  an 
amorphous  substance,  which  is  decomposed  by  water  into  nitric  oxide,  copper,  copper 
nitrite  and  nitrate  [Ber.  26  (1893),  iv,  361]. 

Nitrosylsulphuric  Acid,  SO^NH  — S02<Qf|^^^,  termed  nitrosul- 
phonic  acid,  is  an  intermediate  product  in  the  manufacture  of  commercial 
sulphuric  acid  (see  p.  190),  and  is  quite  important  for  the  determina- 
tion of  the  nitrogen  oxides.  It  is  produced  by  conducting  nitrogen 
tetroxide  into  concentrated  sulphuric  acid  : 

b<^h<CoH  ~ ' -f-  NO2 . OH. 


2o8 


INORGANIC  CHEMISTRY. 


Nitrogen  monoxide,  NO,  is  not  absorl)cd  by  pure  suli)huric  acid,  but 
will  be  if  the  same  contains  nitric  acid  : 


3H,S0.  + HNO3  + 2NO  = 3S03<J]-yo  + 211,0. 

Therefore  it  results  upon  conducting  nitrous  acid  vapors  into  concentrated 
sulphuric  acid : 

2S0,<0[{  + NO,  + NO  = 2S0,<^„^”  + 11,0. 

Further,  the  nitrosylsulphuric  acid  results  from  the  combined  action  of 
sulphur  dioxide,  nitrogen  tetroxide,  oxygen,  and  a little  water: 

2SO,  + N,0.  + 0 + 11,0  = 2S0,<^',^*^. 

It  is  obtained  most  readily  by  conducting  suliihur  dioxide  into  strongly 
cooled  anhydrous  nitric  acid  : 

so,  + UNO,  = 

there  results  a thick  magma,  which  may  be  dried  upon  porous  earthen 
plates  under  the  desiccator. 

Nitrosylsulphuric  acid  forms  a leafy  or  granular  crystalline,  white 
mass  (lead-chamber  crystals,  p.  191),  which  melts  at  about  73°  and  decom- 
poses into  its  anhydride,  sulphuric  acid,  and  nitrogen  oxides.  It  deli- 
quesces in  moist  air,  and  yields  sulphuric  and  nitrous  acids  with  water  : 

+ ^=0  = S0=<0H  + NO.  OH  ; 

the  nitrous  acid  decomposes  further  into  nitric  acid  and  nitric  oxide 
(p.  205). 


Nitrosylsulphuric  acid  dissolves  in  concentrated  sulphuric  acid  without  any  change  ; 
the  solution,  called  nitroso-add,  is  also  produced  in  the  sulphuric  acid  manufacture  in 
the  Gay-Lussac  tower  (p.  190),  is  very  stable  and  may  be  distilled  without  decomposi- 
tion. When  diluted  with  water  it  remains  unaltered  at  first,  but  when  the  specific 
gravity  of  the  solution  reaches  1.55-1.50  (51-48°  Be. ),  then  all  the  nitrogen  oxides  escape, 
especially  on  warming.  When  the  nitroso-acid  is  poured  into  a large  quantity  of  water, 
the  nitrososulphuric  acid  breaks  down  into  sulphuric  and  nitrous  acids,  and  the  latter  in 
part  into  nitric  acid  and  nitric  oxide. 

All  the  nitrogen  oxides  and  acids  are  separated  as  nitric  oxide  (NO)  on  shaking  the 
nitro.so-acid  with  mercury — a procedure  serving  equally  well  for  estimating  the  amount 
of  nitroso-acid  by  means  of  the  nitrometer.  Sulphur  dioxide  acts  similarly  and  upon 
this  depends  the  denitrating  action  of  the  Glover  tower  (see  ]:)p.  189,  190). 

The  anhydride  of  nitrosylsulphuric  acid,  S^N.^Oy  = O | C)  N( )’  P^'^^biced  upon 

heating  the  acid  (together  with  sulphuric  acid,  nitric  oxide,  and  nitrogen  dioxide)  beyond 
its  point  of  fusion.  It  is  obtained  pure  by  .saturating  sulphur  trioxide  with  nitric  oxide  : 

3S().,  -I  2N()  = ()(S(),,.O.NO),,  -t-  SO.,. 


1 1 is  a crystalline;,  colorless  mass,  molting  at  217°,  and  boiling  without  decomposition  at 
about  360°.  Much  water  decomposes  it,  the  same  as  nitrosylsulphuric  acid. 


NITRIC  OXIDE. 


209 


The  chloranhydride  of  nitrosylsulphuric  acid,  SNO^Cl  = S02<Cq‘  is  formed  by  the 

union  of  sulphur  trioxide  with  nitrosyl  chloride  : 

SO3  + NOCl  = SO^C^O  . NO). 

It  forms  white  leaflets,  is  decomposed  by  heat  into  its  components,  and  with  water  breaks 
down  into  sulphuric,  hydrochloric,  and  nitrous  acids. 


Nitric  Oxide,  NO.  When  different  metals  are  dissolved  in  some- 
what dilute  nitric  acid  this  oxide  is  formed.  It  is  conveniently  ob- 
tained by  pouring  dilute  nitric  acid  (specific  gravity  1.2)  upon  copper 
filings : 

3Cu  -f  8HNO3  = 3Cu(N03)2  + 4H2O  -f  2NO. 

A better  procedure  consists  in  pouring  a saturated  solution  of  salt- 
peter upon  thin  copper  turnings  and  then  gradually  introducing  concen- 
trated sulphuric  acid,  or  a concentrated  sodium  nitrite  solution  is 
allowed  to  flow  gradually  into  a hydrochloric  acid  solution  of  ferrous 
chloride  or  sulphate : 

FeCb  + NaNOj  + 2HCI  = FeClg  -f  NaCl  + HjO  + NO. 

The  evolution  begins  in  the  cold  [see  Thiele,  Ann.  Chem.  (1889)  253, 

246]- 

A constant  stream  of  very  pure  nitric  oxide  may  be  obtained  by  pour- 
ing concentrated  sulphuric  acid,  containing  about  2 per  cent,  of  sodium 
nitrite  [Emich,  Monatsh.  f.  Chem.  (1892),  xiii,  73],  upon  mercury. 
Also  by  mixing  a nitrite  solution  with  one  of  potassium  ferrocyanide  and 
acetic  acid  [Deventer,  Ber.  26  (1893),  i,  589]. 

A colorless  gas  escapes,  which,  however,  immediately  forms  brown 
vapors  when  it  comes  in  contact  with  the  air,  as  it  unites  with  the  oxygen 
of  the  latter  to  form  nitrogen  dioxide.  Therefore,  all  the  air  must  be 
expelled  from  the  generating  vessel  by  nitric  oxide,  and  the  gas  collected 
over  water  after  the  interior  of  the  apparatus  has  become  colorless. 

Nitric  oxide  is  a colorless  gas.  Under  normal  conditions  one  liter  of  it 
weighs  T.3426  grams;  its  gas  density  also  remains  unchanged  at — 100° 
(in  comparison  with  air  of  the  same  temperature,  V.  Meyer).  It  is  con- 
densed with  difficulty.  Its  critical  temperature  is  — 93°,  and  its  critical 
pressure  71  atmospheres  (p.  48).  Liquid  nitric  oxide  is  colorless.  It 
boils  at  — 153-6°  and  solidifies  at  — 167°  to  a white  mass.  It  is  slightly 
soluble  in  water,  but  dissolves  very  readily  in  an  aqueous  solution  of  fer- 
rous salts,  imparting  a dark-brown  color  to  the  liquid  ; heat  expels  it 
from  the  same. 

Nitric  oxide  is  readily  soluble  in  nitric  acid.  As  its  solution  becomes  more  concen- 
trated, it  assumes  a brown,  yellow,  green  or  blue  color,  as  nitrogen  trioxide  is  formed 
finally  with  anhydrous  nitric  acid  : 

2IINO3  + 4NO  = 3N2O3  + H2O. 

Potassium  permanganate  oxidizes  it,  like  nitrous  acid  (p.  206),  to  nitric  acid  : 
loNO  4-  6KMnO^  -f  9H2SO*  = 10IINO3  p 3K2SO,  -f  fiMnSO^  + 4H2O. 


210 


INORGANIC  CHEMISTRY. 


As  nitric  acid  contains  53.3  |)cr  cent,  of  oxygen,  it  is  capable  of  sustain- 
ing the  combustion  of  some  substances,  but  only  those  in  whose  com- 
bustion enough  heat  is  liberated  to  effect  the  breaking  down  of  nitric 
oxide  into  nitrogen  and  oxygen.  Hence,  iihosphorus  continues  to  burn 
in  this  gas,  while  a sulphur  hame,  developing  only  a slight  heat,  is  ex- 
tinguished; ignited  charcoal  does  the  same,  while  a sihinter,  burning 
energetically,  will  continue  to  do  so,  when  introduced  into  the  gas. 
On  shaking  a cylinder  filled  with  nitric  oxide  with  a few  drops  of  readily 
volatile  carbon  bisulphide,  and  bringing  a flame  to  the  mouth  of  the 
vessel,  the  carbon  bisulphide  va])ors  will  (piietly  burn  in  the  gas,  giving 
a bright,  luminous  flame,  emitting  strong  actinic  rays ; in  this  combustion, 
the  carbon  and  suliihur  of  the  carbon  bisulphide  unite  with  the  oxygen 
of  the  nitric  oxide  and  form  carbon  dioxide  and  sulphur  dioxide. 

Nitric  oxide  is  a strongly  endothermic  compound  (see  p.  212),  and  is  consequently 
exjdosive  (p.  30). 

On  determining  the  (juantity  of  heat  disengaged  in  the  combustion  of  jdiosphorus,  car- 
bon or  other  sul)stances  in  this  gas,  it  will  be  discovered  that  the  same  is  greater  (about 
21.5  Cal.)  than  that  which  is  developed  by  the  combustion  of  these  bodies  in  oxygen. 
This  can  only  be  explained  upon  the  theory  that  less  heat  is  neces.sary  for  the  separation 
of  NO  into  N and  O than  for  the  separation  of  the  molecules  of  combined  oxygen  atoms 
— an  additional  proof  that  the  molecules  of  free  oxygen  (as  of  other  elements)  consist  of 
atoms. 


With  oxygen  or  air,  nitric  oxide  at  once  forms  brown  vapors  of  nitro- 
gen dioxide  : 


2NO  -b  O2  = 2NO.,. 


2 vols.  I vol.  2 vols. 


With  less  oxygen,  nitrogen  trioxide  is  produced  (p.  205).  Nitric  oxide 
also  combines  with  chlorine  to  nitro.syl  chloride,  NOCl  (p.  204),  and  with 
bromine  it  yields  nitrosyl  bromide,  NOBr.  At  a red  heat  nitric  oxide 
becomes  nitrogen  dioxide  and  nitrogen.  With  hydrogen  and  moderate 
heat  it  forms  water  and  nitrogen  : 

NO  + = N + H.p  ; 

a mixture  of  both  gases  burns  with  a greenish-white  flame.  On  conduct- 
ing a mixture  of  nitric  oxide  and  hydrogen  over  platinum  sponge,  water 
and  ammonia  are  produced  : 

2NO  + 5H2  = 2NH3  + 2H2O. 

The  volumetric  analysis  of  nitric  oxide  gas  may  be  easily  executed  as 
follows:  Fill  a bent  glass  tube  over  mercury  with  the  gas;  introduce 
into  the  same  a piece  of  sodium  and  heat  the  latter  with  a lamp.  The 
sodium  combines  with  the  oxygen,  and  free  nitrogen  separates;  the 
volume  of  the  latter  always  ecpials  half  the  volume  of  the  nitric  oxide 
gas  emi)loyed  ; this  follows  from  the  formula  NO  : 

2NO  = N.2  + O.,. 

2 vols.  I vol.  1 vol. 

Tn  nilrogeti  dioxide  nitrogen  apjicars  to  1)0  ((uadrivalent,  in  nitric  oxide  bivalent,  while 
usually  it  is  Irivalent  or  (|uin(|uivalenl.  Attention  has  already  been  called  to  the  similarity 
of  nitrogen  dioxide  and  chlorine  dioxide,  CIO.^  (p.  207). 


NITROUS  OXIDE. 


21  I 


Nitrous  Oxide,  N.^0,  is  formed  when  zinc  or  tin  acts  ni)on  dilute 
nitric  acid  of  specific  gravity  i.i.  It  may  be  best  obtained  by  heating 
ammonium  nitrate,  which  at  about  170°  breaks  down  directly  into  water 
and  nitrous  oxide : 

NH.NOs  = N,0  + 2H,0. 

It  is  very  probable  that  nitraniide,  NOg . NH2,  is  produced  transitionally,  but  it  immedi- 
ately decomposes  into  water  and  nitrous  oxide  (p.  204).  The  decomposition  is  very 
energetic  at  about  240°  and  is  frequently  accompanied  with  explosion.  It  is  advisable 
to  use  a well-dried  mixture  of  equivalent  quantities  of  sodium  nitrate  and  ammonium  sul- 
phate. The  reaction  will  then  proceed  quietly,  beginning  at  230°,  and  no  explosion  will 
occur  (Smith). 

The  formation  of  the  oxide  from  hydroxylamine  nitrite  (the  isomeride  of  ammonium 
nitrate)  is  interesting.  The  decomposition  takes  place  in  aqueous  solution  and  proceeds 
according  to  the  equation  : 

NO.2  . NH3OH  = N.p  + 2H2O. 

Nitrous  oxide  is  a colorless  gas,  of  sweetish  taste  and  slight  odor.  Its 
density  corresponds  to  the  molecular  formula  N.^O  = 44.08.  One  liter 
weighs  under  normal  conditions  1.9706  grams.  In  cold  water  it  is  toler- 
ably soluble  (i  volume  of  water  dissolves  at  0°  1.305  volumes  of  nitrous 
oxide)  ; therefore  it  must  be  collected  over  warm  water  or  mercury.  By 
pressure  and  cold  it  condenses  to  a colorless  liquid  of  specific  gravity 
0.937  and  boils  under  the  ordinary  pressure  at  — 89.8°.  By  rapid  evapo- 
ration, liquid  nitrous  oxide  solidifies  to  a crystalline  mass  which  remelts 
at — 102.3°.  If  the  aqueous  nitrous  oxide  be  evaporated  over  an  air- 
pump  its  temperature  falls  to  — 140°.  Its  critical  temperature  is  -{-38.8°  ; 
critical  pressure,  775  atmospheres  (p.  48). 

Although  this  oxide  contains  less  oxygen  than  nitric  oxide,  it  supports 
the  combustion  of  many  bodies  more  readily  than  the  latter,  because  it  is 
more  easily  decomposed  into  oxygen  and  nitrogen.  A glimmering  chi}) 
inflames  in  it,  as  in  oxygen  ; phosphorus  burns  in  it  with  a bright,  luminous 
flame,  while  a sulphur  flame  liberating  very  little  heat  is  extinguished.  A 
mixture  of  equal  volumes  of  nitrous  oxide  and  hydrogen  explodes  like 
detonating  gas,  only  less  violently  : 

N2O  -f  H2  = N2  -f  HjO. 

I vol.  I vol.  I vol.  I vol. 

It  resembles  oxygen  very  much,  but  can  be  distinguished  from  it  by 
not  producing  brown  vapors  with  nitric  oxide,  as  does  the  former.  It 
is  not  capable  of  combining  with  oxygen.  When  it  is  conducted 
through  a red-hot  tube  it  is  decomposed  into  nitrogen  and  oxygen.  It 
has  an  exhilarating  effect  when  inhaled  in  slight  quantity,  and  is,  there- 
fore, termed  laughing  gas. 

Its  volume  composition  may  be  determined  in  the  same  manner  as 
with  nitric  oxide,  viz.,  by  heating  a definite  volume  of  the  gas  with 
potassium.  Then  we  learn  that  from  one  volume  of  nitrous  oxide  an 
equal  volume  of  nitrogen  will  be  separated — corresponding  to  the  molec- 
ular formula : 

N./!  + K.2  = N.2  + K2O. 

I vol.  1 vol. 


212 


INORGANIC  CMF.MISTRY. 


Hyponitrous  Acid,  — IIO-N  N-()1I.  In  tlie  list  of  reduction  prodm  Is 

of  nitric  acid  extending  as  far  down  as  ammonia  : 

lINOg,  UNO,,  (UNO),,  Nil, Oil,  NII3, 

this  acid  occupies  a position  between  nitrous  acid  and  hydroxylamine.  Accordingly,  it 
can  be  obtained — 

(1 ) by  the  reduction  of  nitric  and  nitrous  acids  ; 

(2)  by  the  oxidation  of  hydroxylamine  ; and 

(3)  by  the  union  of  nitrous  acid  and  hydroxylamine  with  the  exit  of  water. 

The  reduction  of  the  salts  of  nitric  or  nitrous  acids  may  be  effected  by  sodium  amal- 
gam or  by  electrolysis  ; the  oxidation  of  hydroxylamine  in  acjueous  solution  by  various 
metallic  oxides,  but  especially  by  mercuric  oxide:  2NII2OII  -(-  2llg()  = IKJN 
NOll  (-  2ll,0  T 2llg.  To  obtain  hyponitrous  acid  according  to  (3)  allow  sodium 
nitrite  and  hydroxylamine  sulphate  to  interact,  when  sodium  sulphate  and  hydroxylamine 
nitrite  will  result:  2NaN(),  -j-  SO^(NH^())2  r-  Na,S()^  -f-  2N(J2.  Nil/),  which  for  the 
most  part  breaks  down  into  water  and  nitrous  oxide  (p.  212),  but  in  small  amount  to 
water  and  hyponitrous  acid:  IIONII,  -t-  ON.  OH  — HON  NOH  11,0.  Divers 
claims  that  good  yields  of  potassium  hyponitrite  can  be  obtained  by  boiling  ])otassium 
hydroxylamine  sulphonate  (p.  197)  with  caustic  potash  : 2NH(OH)S03K  -f-  4K(J)H  = 
K2N.,0,  -f-  2lv2'^^\3  411,0.  Upon  neutralizing  the  hyponitrite  solution,  obtained  by 

one  or  the  other  method,  with  acetic  acid  and  adding  silver  nitrate  the  silver  salt  Ag.^N,!), 
separates  as  a bright  yellow,  amorphous  powder,  which  slowly  decomposes  at  100°,  but  at 
110°  with  violent  explosion.  Concentrated  sulphuric  acid  liberates  nitrous  oxide  from  it. 

The  acid  may  be  obtained  in  ethereal  solution  by  adding  gradually  an  excess  of  its 
silver  salt  to  a solution  of  dry  hydrogen  chloride  in  anhydrous  ether.  After  evaporating 
the  ethereal  solution  in  a vacuum,  free  hyponitrous  acid  remains  in  crystalline  leaflets.  It 
explodes  readily — with  flame — oncoming  in  contact  with  solid  caustic  potash  ; deliquesces 
in  the  air,  is  very  soluble  in  alcohol,  and  soluble  in  benzene  and  in  chloroform.  It  is 
more  stable  in  aqueous  solution  than  in  the  solid  condition,  but  even  then  readily  decom- 
poses into  nitrous  oxide  and  water:  HgN.^O,  = N,0  -)-  11,0.  It  therefore  bears  the 
same  relation  to  nitrous  oxide  as  nitric  acid  to  nitrogen  pentoxide,  and  nitrous  acid  to 
nitrogen  trioxide  ; however,  nitrous  oxide  does  not  take  up  water  and  it,  therefore,  does 
not  behave  like  the  true  anhydride  of  hyponitrous  acid.  Upon  neutralizing  its  aqueous 
solution  with  caustic  soda  the  acid  salt  NaHN202  is  produced,  which  quickly  breaks 
down  into  caustic  soda  and  nitrous  oxide  : NaHN.,^^,  = N,!^  + NaOH.  The  freshly 
prepared  solution  of  the  acid  liberates  iodine  from  potassium  iodide  after  it  has  stood  for 
some  time  ; this  is  because  it  decomposes  very  slowly  and  only  to  a slight  degree  into 
nitrous  acid  and  ammonia:  3H2N2O2  =:  2N2O3  -j-  2NH3.  It  decolorizes  potassium 
permanganate  and  becomes  nitric  acid. 

The  molecular  magnitude  of  the  acid  was  established  by  the  determination  of  the  vapor 
density  of  its  ethyl  ester,  N202(C2H5)2,  by  Zorn,  and  from  the  molecular  weight  of  the 
benzyl  ester,  N202(C7lI.j)2,  by  Hantszch  and  Kaufmann. 

Hyponitrous  acid  is  isomeric  with  nitramide  (p.  204). 

Compare  the  investigations  of  Hantzsch  and  his  co-workers  : Ann.  Chem.  292  (1896), 
317  ; 299  (1898),  94;  Divers  ibid.  295  (1897),  366;  also  Kirschner,  Z.  f.  anorg.  Chem. 
16  (1898),  424. 


d'he  thermo-cheinical  relations  of  the  oxygen  derivatives  of  nitrogen  give  some  clue 
to  their  chemical  deportment.  All  nitrogen  oxides  are  endothermic  compounds,  i.  e., 
they  are  produced  from  their  elements  with  heat  absorption  corresponding  to  the  symbols  : 

(N„0)=-i7.5  (N,  0)  = -2i.6  (N,  0,)=-S. 

I'rom  this  we  observe  that  the  oxides  of  nitrogen  cannot  be  ]'>repared  from  the  elements 
without  addition  of  energy.  Proceeding  from  nitric  oxide  (Nt)),  we  see  from  the 
above  numbers  that  the  formation  of  the  liigher  oxides  from  it  occurs  with  heat  disen- 
gagetnent  : 


(2NO,  i))  ■=  20.1  (NO,  O)  = 13.4  (2NO,,  (),)  = 40.5, 


OXYGEN  COMPOUNDS  OF  PHOSPHORUS. 


213 


whereas  lieat  is  absorbed  in  the  conversion  of  nitrous  into  nitric  oxide  : (N2O,  O)  = 

— 25.4.  (See  J.  'I'homsen’s  Thermocheniische  Untersuchungen,  iv,  314.) 

lieat  disengagement,  on  the  contrary,  occurs  in  the  production  of  nitric  acid  from  its 
elements : 

(N,03,H-liquid)  =41.6  (N,03,H,Aq)  = 49.1. 

This  explains  the  relative  stability  of  that  acid. 


Compounds  of  Nitrogen  with  Sulphur. — Two  such  bodies  are  known  : nitrogen 
sulphide,  N^.S^,  corresponding  in  composition  to  a quadrupled  nitric  oxide,  and  nitrogen 
pentasulpJnde,  corresponding  to  nitric  anhydride,  NgO^. 

Nitrogen  sulphide,  N4S4,  is  produced  when  dry  ammonia,  in  benzene  solution,  acts 
upon  sulphur  dichloride  : 

4NH3  + 6SCI2  =:  N4S4  + I2HCI  + 83. 

It  consists  of  orange-red  needles,  melting  at  178°.  It  detonates  at  higher  temperatures 
or  when  struck.  Its  molecular  weight  has  been  determined  by  methods  which  will  be 
described  under  Solutions. 

Nitrogen  pentasidphide,  N4S5,  is  formed  by  a complicated  reaction  occurring  on  heat- 
ing nitrogen  sulphide  with  carbon  bisulphide  to  100°  in  a sealed  tube.  It  is  a deep-red 
liquid,  transparent  in  thin  layers  with  a blood-red  color.  It  is  very  mobile,  not  moisten- 
ing  glass  and  its  specific  gravity  at  18°  equals  1.90.  It  solidifies  by  cold  to  a mass  resem- 
bling iodine  and  melts  at  10°.  Its  odor  is  like  that  of  iodine  and  is  at  the  same  time 
sw'eet.  Nitrogen  pentasulphide  is  not  soluble  in  water,  but  in  organic  solvents.  Its  solu- 
tions are  more  stable  than  the  pure  pentasulphide,  w Inch  easily  decomposes  into  its  ele- 
ments. Boiling  water,  alkalies  and  hydrogen  sulphide  decompose  the  pentasulphide 
very  readily  with  the  production  of  ammonia  and  the  separation  of  sulphur.  [See  Z.  f. 
anorg.  Chem.  13  (1897),  200.] 


2.  OXYGEN  COMPOUNDS  OF  PHOSPHORUS. 


P4O6 

Phosphorus 

trioxide. 

P2O5 

Phosphorus 

pentoxide. 


Hypophosp 

acid. 


horous 


H3PO3 

Phosphorous 

acid. 


Orthophosphoric 

acid. 


Oxides  containing  less  oxygen,  e.  g.,  the  tetroxide  and  the  oxide 
P./j,  are  not  definitely  known. 

The  following  anhydride  acids  are  derived  from  orthophosphoric  acid  : 


HPO3,  Metapho.sphoric  acid. 

H4P2O7,  Pyrophosphoric  acid. 

Orthophosphoric  acid  also  yields  an  anhydride  acid  with  phosphorous 
acid  : hypophosphoric  acid,  H4P20g. 

The  structure  of  these  compounds  is  expressed  by  the  following 
formulas : 

V HI  V OH  V 

II3PO-OII  p ^OII  iikxCh 

Hypophosphorous  Phosphorous  Phosphoric 

acid.  acid.  acid. 


214 


INORGANIC  CHEMISTRY. 


In  hypophos])horoiis  acid  two  atoms  of  liydrogcn  are  in  direct  union 
with  quinquivalent  phospliorus,  while  the  third  hydrogen  atom  forms  an 
hydroxyl  group  with  oxygen.  It  is  only  this  last  atom  of  hydrogen  which 
is  readily  replaced  by  the  action  of  bases ; hence  hypophosphorous  acid 
is  a 7nonobasic  acid.  Phosphorous  acid  (like  sulphurous  and  arsenious 
acids)  appears  in  two  forms  in  its  derivatives  (pp.  i86,  221).  It  is  very 
probable  that  in  its  salts  it  has  one  atom  of  hydrogen  and  two  hydroxyl 
groups  joined  to  ])hosphorus;  consecjuently  it  is  a dibasic  acid.  [See 
Michaelis  and  Pecker,  Per.  30  ( 1897),  1003.]  Finally,  three  series  of  salts 
and  three  hydroxyls  are  assumed  to  be  present  in  phosphoric  acid.  Py 
the  elimination  of  one  molecule  of  water  from  ])hosi)horic  acid,  meta- 
])hosphoric  acid  results — an  anhydride  which,  at  the  same  time,  is  a 
monobasic  acid,  as  it  contains  one  hydroxyl  group  : 


V 

POj-OII,  Metaphosphoric  acid. 

On  removing  one  molecule  of  water  from  two  molecules  of  jihosphoric 
acid,  pyro-  or  diphosphoric  acid  is  formed  (see  }).  193)  : 


V /OH 
PO^OH 

^OH 

V /OH 
PO^OH 

\OH 


— H,0 


2 molecules  Phosphoric  acid. 


/OH 

PO^OH 

PO— OH 
^OH. 


I molecule  Pyrophosphoric  acid. 


Pyrophosphoric  acid  contains  four  hydroxyl  groups,  hence  is  tetra- 
basic.  Similarly,  hypophosphoric  acid  results  from  one  molecule  each 
of  ])hosphoric  acid  and  phosphorous  acid,  P(OH)3. 

Finally,  if  from  two  molecules  of  phosphoric  acid  all  the  hydrogen  atoms 
be  removed,  in  the  form  of  water,  an  anhydride  remains : 


V V 
O2P-O-PO2. 

Phosphoric  anhydride. 

The  salts  of  phosphoric  acid  are  termed  phosphates ; those  of  phos- 
phorous acid,  phosphites^  and  of  hypophosphorous  acid,  hypophosphites. 


Hypophosphorous  Acid,  H3PO2.  Hydrogen  phosphide  e.scapes 
when  a concentrated  solution  of  sodium  or  barium  hydroxide  is  warmed 
with  yellow  phosphorus,  leaving  behind  in  solution  a salt  of  hypophos- 
phorous acid  : 

4?  4-  3NaOII  -p  3II.P  = 3TT.TO.ONa  + PH, 

8P  -f  3l5a(OHh  + 3II2O  = 3(Il2l'0. 0)Ta+  2Pli3. 

d'he  free  acid  may  be  separated  from  the  barium  salt  by  means  of  sul- 
])huric  acid  ; the  insoluble  barium  suljihate  being  filtered  off  from  the 
acpieous  solution  of  the  acid,  and  the  latter  concentrated  under  the  air- 


PHOSPHOROUS  ACID. 


215 


pump.  Hypophosphorous  acid  is  a colorless,  thick  liquid,  with  a strong 
acid  reaction.  Below  0°  it  sometimes  solidifies  to  large,  white  leaflets, 
which  melt  at  Heat  converts  it,  with  much  foaming,  into 

hydrogen  phosphide  and  phosphoric  acid  : 

2H3P0,^PH3  + II3PO,. 

It  absorbs  oxygen  readily,  becoming  phosphoric  acid,  hence  acts  as  a 
powerful  reducing  agent.  It  reduces  sulphuric  acid  to  sulphur  dioxide, 
and  even  to  sulphur.  It  precipitates  many  of  the  metals  from  their  solu- 
tions ; from  copper  sulphate  it  separates  the  hydride  CiqH^. 

The  acid  is  monobasic,  H2PO.OH  (see  p.  214).  Its  salts  dissolve 
readily  in  water,  and  absorb  oxygen  from  the  air,  thus  becoming  phos- 
phates. When  heated  in  a dry  condition,  they  set  free  phosphine  and 
hydrogen  and  are  converted  into  pyro- and  metaphosphates;  some  also 
yield  metallic  phosphides.  Nascent  hydrogen  reduces  hypophosphorous 
acid  to  phosphine. 


Phosphorous  Acid,  H3PO3,  is  formed  at  the  same  time  with  phos- 
phoric acid  and  hypophosphoric  acid  in  the  slow  oxidation  of  phosphorus 
in  moist  air.  The  decomposition  of  phosphorus  trichloride  by  water  gives 
it  more  conveniently : 

PCI3  + 3H2O  = H3PO3  + 3HCI. 

By  evaporating  this  solution  under  the  air-pump  the  phosphorous  acid 
becomes  crystalline.  The  crystals  are  readily  soluble  in  water,  and 
deliquesce  in  the  air.  It  melts  at  70°,  and  decomposes  on  further  heating 
into  phosphine  and  phosphoric  acid  : 

4H3PO3  = PH3  + 3H3PO,. 

In  the  air  the  acid  absorbs  oxygen,  and  changes  to  phosphoric  acid. 
Hence,  it  is  a strong  reducing  agent,  and  precipitates  the  free  metals 
from  many  of  their  solutions.  In  the  presence  of  water  the  halogens 
oxidize  it  to  phosphoric  acid. 

It  is  a dibasic  acid,  forming  two  series  of  salts,  in  which  one  or  two 
atoms*  of  hydrogen  are  replaced  by  metals.  The  phosphites  do  not  oxi- 
dize in  the  air,  except  under  the  influence  of  strong  oxidizing  agents 
When  heated,  they  generally  decompose  into  hydrogen,  pyrophosphates 
and  phosj)hide.  Nascent  hydrogen  also  reduces  phosphorous  acid  to  phos- 
phine— a circumstance  of  importance  in  the  detection  of  phosphorus- 
poisoning. 

Phosphorous  Anhydride,  or  phosphorus  trioxide  according  to  the  old  formula  P2O3, 
has  been  shown  by  4'horpe  and  Tutten  (1892)  to  be  correctly  expressed  by  1^)6-  B is 


* Therefore,  the  structural  formula,  TIPO(OII)2  is  assigned  to  this  acid.  There 
appears  to  exist  another  phosphorous  acid,  at  least  in  compounds,  to  which  the  formula 
P(OII)3  belongs  (pp.  213,  216). 


2i6 


INORGANIC  CUKMISTRV. 


produced  on  conducting  dry  air  over  gently  lieated  jdiospliorus,  or  by  carefully  beating 
phosphorous  acid  with  i)hosphorus  trichloride  : 

2II3PO3  + 2PCI3  =:  P,Og  + 6IIC1. 

It  is  a white,  flocculent  mass  or  it  consists  of  colorless  needles.  It  melts  at  22.5°,  sub- 
limes readily  and  boils  in  an  atmosphere  of  nitrogen  at  173°.  I ts  vapor  density  corre- 
sponds to  the  formula  I^Og.  At  400°  it  breaks  down  into  phosphorus  and  phosphorus 
tetroxide,  PjO^,  which  crystallizes.  It  is  decomposed  by  water  in  a very  complicated 
manner. 


Phosphoric  Acid,  H3P0^,  or  Orthophosphoric  acid,  is  produced  when 
the  peiitoxide  is  dissolved  in  hot  water,  and  by  the  decomposition  of 
phosphorus  pentachloride  or  phosphorus  oxychloride  (POCl,)  by  water  (see 
p.  219).  It  may  be  obtained  by  decomposing  bone  ash,  Ca3(PO^)2,  with 
sulphuric  acid,  or  better,  by  oxidizing  yellow  phosithorus  with  nitric  acid. 
The  aqueous  solution  is  evaporated  to  dryness  in  a platinum  dish.  The 
anhydrous  acid  consists  of  colorless,  hard,  prismatic  crystals,  which  in 
the  air  deliquesce  to  a thick,  acid  liquid. 

Phosphoric  acid  is  tribasic,  forming  three  series  of  salts,  q.2\\^(\ primary 
(KH2P0^),  secondary  (K^HPO^),  and  tertiary  (K3POJ.  I'hey  may  be 
spoken  of  according  to  the  number  of  hydrogen  atoms  rei)laced  by  metals, 
as,  e.  g.,  monopotassium  phosphate  (KH2POJ,  di])otassium  phosphate 
(K2HPO4),  and  tripotassiiim  phosphate  (K3POJ.  The  salts  of  the  first 
two  series  contain  hydrogen,  replaceable  by  metals,  hence  may  be  termed 
acid,  while  the  salts  of  the  third  series  are  neutral.  Their  behavior  with 
litmus  does  not  harmonize  with  this  view  (see  Sodium  Phosphate). 

The  tertiary  phosphates,  excepting  the  salts  of  the  alkalies,  are  insolu- 
ble in  water.  With  a silver  nitrate  (AgNOg)  solution,  soluble  phosphates 
give  diyellow  precipitate  of  silver  phosphate,  AggPO^. 

Pyrophosphoric  Acid,  PI^P207  (structure,  p.  214),  is  formed  by 
the  continuous  heating  of  orthophosphoric  acid  to  260°,  until  a portion 
of  it  dissolved  in  ammonium  hydroxide  does  not  yield  a yellow  but  a pure 
white  preci])itate  with  silver  nitrate.  The  sodium  salt  is  easily  obtained 
by  heating  disodium  phosphate: 

2Na2HPO^  = NaTjO;  + H2O. 

Other  salts  are  similarly  formed  by  heating  the  corresponding  ortho- 
phosphates. 

'Fhe  acid  presents  a white  crystalline  appearance,  and  is  readily  soluble 
in  water.  When  in  solution,  it  slowly  takes  up  water  at  ordinary  tem- 
])eratures,  more  rapidly  when  heated,  and,  like  all  true  anhydrides,  passes 
into  the  corresponding  acid — orthophosi)horic  acid. 

Pyrophosphoric  acid  is  tetrabasic.  Its  salts  are  very  stable,  and  are 
not  altered  by  boiling  with  water  ; warmed  with  dilute  acids,  they  become 
salts  of  the  ortho-acid,  'bhe  soluble  salts  yield  a 7vhife  ])recipitate  of  silver 
])yro]jhosphate,  AgJ^O^,  with  silver  nitrate. 

Hypophosphoric  Acid,  II,P20^^.  While  i^yrophosphoric  acid  is  an  anliydride  acid 
of  |)hoH])lioric  acid,  the  so-called  hy|)ophosplu)ric  may  he  viewed  as  a mixed  anhydride 
of  pliosj)lioric  and  symmelrical  phosj)horoiis  acids. 


METAPHOSPHORIC  ACID. 


217 


PO(OH), 


It  is  produced,  as  demonstrated  by  Salzer  in  1877,  together  with  phosphorous  and  phos- 
phoric acids,  by  the  slow  oxidation  of  moist  phosphorus  in  the  air.  It  is  separated  from 
these  acids  by  means  of  its  difficultly  soluble  sodium  salt,  NajH.^P.^Og  -(-  6H.^O  ; by  pre- 
cipitating the  solution  of  the  latter  with  a soluble  lead  salt  we  get  insoluble  lead  hypo- 
phosphate,  PbjP.^Og.  Its  silver  salt  is  more  easily  obtained  by  oxidizing  phosphorus  in 
the  presence  of  silver  nitrate.  The  free  acid  separated  from  the  lead  or  silver  salt  by 
hydrogen  sulphide  is  rather  stable  in  a dilute  solution,  and  below  30°  may  be  concen- 
trated to  a syrup.  At  higher  temperatures,  more  readily  in  the  presence  of  hydrochloric 
or  sulphuric  acid,  the  acid  decomposes  into  phosphoric  and  phosphorous  acids.  It  does 
not  reduce  metallic  salts,  but  is  oxidized  by  potassium  permanganate  to  phosphoric  acid. 
[Salzer,  Ann.  Chem.  211  (1892),  i.  See  Z.  f.  anorg.  Chem.  6 (1894),  132,  for  a con- 
venient method  for  its  production.] 

Metaphosphoric  Acid,  HPO3  or  PO2 . OH,  results  upon  heating  the 
ortho-  or  pyro-acid  to  300°.  It  can  be  more  conveniently  obtained  by 
dissolving  phosphorus  pentoxide  in  cold  water : 


P2O5  -f  HjO  = 2HPO3. 


It  is  a vitreous,  transparent  mass  (^Acidum pJwsphoricum  glaciale),  which 
generally  contains  less  hydrogen  than  is  required  by  the  formula  HPO3. 
This  probably  is  because  some  anhydride  is  present  in  it.  It  melts  when 
heated  and  volatilizes  at  higher  temperatures,  without  suffering  any 
change.  It  deliquesces  in  the  air,  and  dissolves  with  ease  in  water. 
(Commercial  glacial  phosphoric  acid  contains  sodium  and  magnesium 
phosphate,  and  dissolves  with  difficulty  in  water.)  The  solution  coagu- 
lates albumin ; this  is  a characteristic  method  of  distinguishing  meta- 
from  ortho-  and  pyrophosphoric  acid.  In  aqueous  solution,  metaphos- 
phoric acid  changes,  gradually  at  ordinary  temperature,  rapidly  by  boil- 
ing, into  orthophosphoric  acid  : 


HPO3  + H2O  = H3PO,. 


It  is  a monobasic  acid.  Its  salts,  the  metaphosphates,  are  readily 
obtained  by  the  ignition  of  the  primary  salts  of  the  ortho-acid : 


Nall^PO,  = NaP03  -f-  H^O. 


When  the  aqueoussolutions  of  these  salts  are  boiled,  they  are  converted 
into  the  primary  salts  of  ortho})hosphoric  acid.  With  silver  nitrate  the 
soluble  metaphosphaies  give  a white  precipitate  of  silver  metaphosphate. 


In  addition  to  the  ordinary  salts  of  metaphosphoric  acid,  various  modifications  of  the 
same  exist ; these  are  derived  from  the  polymeric  meta-acids,  ri2P.20g,  H3P3O,,,  H^P^O,2, 
etc.  They  are  all  changed  to  primary  ortho-phosphates  by  boiling  their  solutions.  [See 
Tammann,  Jr.  prakt.  Ch.  45  (1892),  417.] 

Phosphorus  Pentoxide,  PaO^,  or  Phosphoric  anhydricte,  is  obtained 
by  burning  phosphorus  in  a current  of  dry  oxygen  or  dry  air. 

19 


2i8 


INORGANIC  CHEMISTRY. 


The  following  procedure  serves  for  the  jjrcparation  of  it  (Fig.  65)  : A piece  of  phos- 
phorus, placed  in  an  iron  dish  attached  to  the  glass  tube  a b,  is  burned  in  the  glass 
balloon  A.  The  necessary  amount  of  air  is  drawn  through  the  vessel  by  means  of  an 
aspirator.  The  air  is  lirst  ])assed  through  the  bent  tube  containing  pieces  of  pumice- 
stone,  moistened  with  sulphuric  acid,  in  order  to  dry  it  perfectly.  After  the  j)hosphorus 
has  been  consumed,  fresh  pieces  of  it  are  introduced  into  the  little  di.sh  through  a /f>,  and 
the  upper  end  of  the  tube  closed  with  a cork.  The  phosphorus  pentoxide  formed  collects 
partly  in  A and  partly  in  the  receiver  (//). 

Phosphorus  pentoxide  is  a wliite,  voluminous,  flocculent  mass.  It 
attracts  moisture  energetically  and  delitpiesces  in  the  air.  It  dissolves  in 
cold  water  with  hissing  and  yields  inetajihosphoric  acid.  Owing  to  its 
great  affinity  for  water  it  serves  as  an  agent  for  drying  gases,  and  also  for 
the  withdrawal  of  water  in  chemical  reactions,  e.  g.,  in  the  formation  of 
acid  anhydrides  from  acids  (p.  204). 

Tt  generally  contains  oxides  of  lower  oxygen  content.  It  can  be  freed  from  these  by 
subliming  it  over  ignited  platinum  sponge  in  a current  of  oxygen. 


Fig.  65. 


Chloranhydrides  of  the  Acids  of  Phosphorus. — The  halogen  derivatives  of  phos- 
phorus, considered  on  p.  141,  may  be  viewed  as  the  halogen  anhydrides  of  phos- 
phorous and  phosphoric  acids  (p.  195).  The  compounds  PCI3,  PBrg,  and  PI3  are 
derived  from  phos})horous  acid,  because  they  yield  the  latter  acid  with  water : 

PCI,  h 311,0  = II3PO3  + 3IICI. 

The  compounds,  jdiosphorus  pentachloride,  PCI5,  and  phosphorus  pentabromide, 
Plhj,  correspond  to  the  normal  orthophosphoric  acid,  P(01I)5,  which  has  not  been 
obtained  in  a free  condition  (.see  Calcium  Phosphate). 

Phosphorus  Oxychloride,  POCI3,  can  be  viewed  as  the  chloride 
of  the  known’ orthophosjihoric  acid,  PC)(()H)3.  It  is  a colorless  liquid, 
fuming  strongly  in  the  air;  its  siiecific  gravity  equals  1.68  at  15°.  It 


COMPOUNDS  OF  PHOSPHORUS  WITH  SULPHUR. 


219 


congeals  on  cooling  and  then  melts  at  — 1-5°.  It  boils  without, decom- 
position at  107°;  its  vapor  density  corresponds  to  the  formula — 

POCI3  = 153.0. 

Water  decomposes  it  intometa-  or  orthophosphoric  acid  and  hydrochloric 
acid : 

rocij  + 2H2O  = H PO3  + 3HCI 
POCI3  + 31102  = HgPO^  + 311  Cl. 

It  may  be  obtained  by  decomposing  the  pentaehloride  with  a little  water 
(see  above),  or  by  allowing  it  to  deliquesce  gradually  in  moist  air.  A 
more  practical  method  consists  in  distilling  the  pentaehloride  with  phos- 
phorus pentoxide  : 

3PCl3  + P203  = 5POCl3; 

or  with  crystallized  boric  acid  (6  parts  to  i part)  : 

3PCI5  -f  2H3BO3  = 3POCI3  + B2O3  + 6HC1. 

A noteworthy  formation  of  the  oxychloride  is  the  union  of  phosphorus 
trichloride  with  oxygen  on  passing  ozonized  air  through  it : 

PCI3  + 03  = POCI3  + O2. 

Potassium  chlorate  acts  in  a corresponding  manner  very  energetically 
upon  the  trichloride  with  the  production  of  oxychloride: 

3PCI3  + KCIO3  = 3POCI3  4 KCl. 

The  compound  PSCI3  is  analogous  to  the  oxychloride  POCI3.  It  is  obtained  by  heat- 
ing the  trichloride  and  sulphur  to  130°  ; also  by  the  action  of  phosphorus  pentaehloride 
upon  hydrogen  sulphide  or  some  metallic  sulphides  : 

PCI5  + H2S  = PSCI3  -f  2HCI. 

Phosphorus  sulphochloride,  PSCI3,  is  a colorless  liquid  of  specific  gravity  1.6,  fuming  in 
the  air  and  boiling  at  124-125°.  Water  decomposes  it  into  phosphoric  and  hydrochloric 
acids  and  hydrogen  sulphide  : 

PSCI3  -f  4H2O  = H3PO4  p 3HCI  + H2S. 

The  chlorides  PO2CI  and  P203C1^,  corresponding  to  meta-  and  pyro- 
phosphoric  acids,  have  also  been  prepared.  (See  Chem.  Centralblatt, 
1897,  II,  I4-) 


COMPOUNDS  OF  PHOSPHORUS  WITH  SULPHUR. 

With  sulphur,  phosphorus  affords  a number  of  compounds  which  are 
obtained  by  direct  fusion  of  phosphorus  with  sulphur.  As  the  union 
of  ordinary  phosphorus  and  sulphur  usually  occurs  with  violent  explo- 
sion, red  phosphorus  should  be  employed  in  preparing  these  compounds. 

The  compounds  PjS,  and  PjS^,  analogous  in  constitution  to  P^O^  and 
PjOj,  are  solid  crystalline  substances,  melting  at  higher  temperatures  and 


220 


INORGANIC  CHEMISTRY. 


subliming  without  decomposition;  phosphorus  pentasulphide  boils  at 
520°  and  the  trisulphide  at  540°  (cor.).  Water  changes  them  to  hydro- 
gen sulphide  and  the  corresponding  acids,  i)hosphorous  and  phosphoric. 
They  combine  with  the  alkaline  suli)hides  to  form  compounds  (e.  g., 
K3PSJ  which  possess  a constitution  analogous  to  that  of  the  salts  of  phos- 
phoric and  phosphorous  acids  [see  Suli)ho-salts  of  Arsenic,  j).  223;  also 
Glatzel,  Z.  f.  anorg.  Chem.  4 (1893),  187]. 

Well-crystallized  sulphur  phosphides,  the  composition  of  which  corre- 
sponds to  the  formulas  P^Sg  and  PgS^,  have  been  i)repared.  The 

supposed  compounds  P^S  and  P^S  are  licpiids  which  inflame  readily  in 
contact  with  air,  but  have  been  proved  to  be  mixtures  containing  free 
phosphorus. 

Besides  the  preceding,  we  have  other  phosphorus  derivatives  which  contain  nitrogen. 
These  have  been  little  studied,  and  at  present  are  unimportant.  Such  compounds 
are  PNjfl  (phospham),  PNO,  PNCI.^.  The  so-called  amid-derivatives,  P()Cl2.Nll2, 
P0C1(NH2)2  and  PO(NH2)s»  are  produced  by  allowing  ammonia  to  act  upon  phosphorus 
oxychloride,  POCI3.  In  these  chlorine  is  replaced  by  the  amido-group  Nlf^. 

Recently,  amidophosphoric  acid,  P0(NIl2) (011)2,  corresponding  to  sulphamic  acid, 

S02(NH2)0H,  and  imidophosphoric  acid,  the  counterpart  of  imido- 

sulphuric  acid,  as  well  as  its  derivatives,  have  been  more  exhaustively  studied  [Stokes, 
Jahrb.  d.  Chem.  vi  (1896),  89]. 


3.  OXYGEN  DERIVATIVES  OF  ARSENIC. 

AS2O3  .... 

Arsenic  trioxide. 

AS2O5  HjAsO^. 

Arenic  pentoxide.  Arsenic  acid. 

Arsenic  Trioxide,  or  Arsenioiis  anhydride  (^Acidum  arseni- 

osum),  occurs  in  nature  as  arsenic  “bloom.”  It  is  produced  by  the 
burning  of  arsenic  in  oxygen  or  in  the  air,  and  by  the  oxidation  of  the 
metal  with  dilute  nitric  acid.  It  is  obtained  metallurgically  on  a large 
scale  as  a by-product  in  the  roasting  of  ores  containing  arsenic.  The 
trioxide  thus  formed  volatilizes  and  is  collected  in  walled  chambers,  in 
which  it  condenses  in  the  form  of  a white  powder  {white  arsenic,  poison 
flour).  To  render  it  pure,  it  is  again  sublimed  in  iron  cylinders,  and 
obtained  in  the  form  of  a transparent,  amorphous,  glassy  mass  {arsenic 
glass),  the  specific  gravity  of  which  equals  3.74.  Upon  preservation  this 
variety  gradually  becomes  opaque  and  porcelaneoiis,  acquires  a crystalline 
structure,  and  its  specific  gravity  decreases  to  3.69.  Upon  dissolving 
this  oxide  in  hot  hydrochloric  acid,  it  crystallizes,  on  cooling,  in  shin- 
ing, regular  octahedra.  At  the  same  time,  this  interesting  phenomenon  is 
observed  : that  when  the  solution  of  the  glass  variety  crystallizes  it  phos- 
phoresces strongly  in  the  dark,  while  the  porcelaneoiis  does  not  exhibit 
this  jiroperty.  Arsenic  trioxide  crystallizes  in  similar  forms  of  the  regu- 
lar system  when  its  vapors  are  rajiidly  cooled,  but  upon  cooling  slowly, 
it  assumes  the  shape  of  monoclinic  jirisms  of  sjiecific  gravity  4.0  ; 
therefore  it  is  dimorphous.  When  heated  in  the  air,  it  sublimes  above 


OXYGEN  DERIVATIVES  OF  ARSENIC. 


221 


218°,  without  melting;  under  higher  pressure,  however  (in  sealed  tubes), 
it  melts  to  a liquid  which  solidifies  to  a glassy  mass. 

The  molecular  magnitude  of  the  solid  arsenic  trioxide,  like  that  of  all  solids,  is  not 
known.  When  converted  into  vapor  it  behaves  like  sulphur  and  other  solid  liquid  bodies  ; 
in  vapor  form  at  lower  temperatures  the  molecules  contain  more  atoms  than  at  higher 
temperatures.  According  to  Biltz  (Z.  f.  phys.  Ch.  20  (1896),  68),  at  temperatures  rang- 
ing from  500°  to  700°  the  molecules  of  As^Og  predominate,  while  between  700°  and  1800° 
the  molecules  of  AsjOg  increase  in  number.  Above  1800°  the  density  corresponds  to  the 
old  formula  AsjO,,  and  for  this  reason  and  simplicity’s  sake  it  is  retained  in  the  text.  It 
is  interesting  to  observe  that  at  lower  temperatures  the  same  group  As^  is  present  in  the 
molecule  of  arsenious  oxide  as  was  observed  in  the  molecule  of  arsenic  (p.  143). 

The  trioxide  dissolves  with  difficulty  in  water  ; the  solution  possesses  a 
sweetish,  unpleasant  metallic  taste,  exhibits  but  feeble  acid  reaction,  and 
is  extremely  poisonous.  The  oxide  is  very  soluble  in  acids,  and  probably 
forms  salts  with  them;  at  least,  on  boiling,  a solution  of  the  trioxide 
volatilizes  in  strong  hydrochloric  acid,  arsenious  chloride,  AsClg.  From 
this  and  its  feeble  acid  nature  we  perceive  an  indication  of  the  basic  char- 
acter of  the  trioxide  corresponding  to  the  already  partly  metallic  nature 
of  arsenic  (see  p.  148). 

Nascent  hydrogen  converts  the  trioxide  into  arsine  (AsHg)  ; but  when 
heated  with  charcoal  it  is  reduced  to  the  metallic  state  (pp.  143,  144). 
Upon  heating  arsenic  trioxide  in  a narrow  glass  tube  with  carbon,  the 
reduced  arsenic  deposits  as  a metallic  mirror  on  the  sides.  Oxidizing 
agents  convert  arsenic  trioxide  into  arsenic  acid. 

Arsenious  Acid,  HgAsOg,  corresponding  to  the  anhydride,  is  not 
known  in  a free  condition.  It  probably  exists  in  the  aqueous  solution 
of  arsenic  trioxide,  but  the  anhydride  separates  out  upon  evaporation. 
In  its  salts  {arsenites)  it  is  tribasic  and  usually  affords  tertiary  derivatives : 

AgsAsOg,  Mg3(AsOg).^. 

The  alkali  salts,  soluble  in  water,  absorb  oxygen  from  the  air  and  serve 
as  powerful  reducing  agents,  they  themselyes  becoming  arseniates.  Silver 
nitrate  forms  a yellow-colored  precipitate,  AggAsOg,  with  the  soluble  salts. 
[See  Stavenhagen,  Jr.  f.  prakt.  Chem.  51  (1894),  i]. 

Other  salts  exist  which  are  derived  from  the  meta-arsenious  acid, 
HASO2  (p.  201). 


According  to  Klinger,  arsenious  acid,  like  sulphurous  and  phosphorous  acids,  is  pres- 

yOW 

ent  in  its  derivatives  in  two  forms : (i)  the  symmetrical  As^OH,  from  which  the  silver 

^OH 


OH 


salt,  the  chloride  and  the  esters  are  derived,  and  (2)  the  unsymmetrical  H-As<q 
or  H - As^OH,  from  which  the  alkali  salts  and  alkyl  arsonic  acids  (see  Organic  Chem- 


^01 1 
istry)  have  their  origin. 


Arsenic  Acid,  HgAsO^  ( Acidtim  arsenicuni),  is  obtained  by  the  oxi- 
dation of  arsenic  or  its  trioxide  with  concentrated  nitric  acid  or  by 
means  of  chlorine.  Upon  evaporating  the  solution  rhombic  crystals  of 


222 


INORGANIC  CHEMISTRY. 


the  forniiila  TI^AsO^  -f-  sej)aralc  out;  these  deliiiuesce  on  exposure 

to  the  air.  'I'liey  melt  at  ioo°,  lose  their  water  of  crystallization  and 
yield  ortho-arsenic  acid,  H3AsO^,  which  heated  to  140-180°  passes  into 
Pyro-arse?iic  acid,  H^As207 : 

2II3ASO,  — . H.AsjO^  -f  II2O. 

At  200°  this  again  loses  water  and  becomes  Meta-arscnic  acid,  HAsOg. 
With  water  the  last  two  acids  become  ortho-  again  ; hence  the  latter  is 
perfectly  analogous  to  idiosphoric  acid.  A portion  of  the  arsenic  acid 
is  volatilized,  i)robably  as  arsenious  chloride,  when  strong  hydrochloric 
acid  solutions  of  arsenic  acid  are  heated. 

At  a red  heat  the  meta-arsenic  acid  loses  all  its  water  and  becomes 
Arsenic  pentoxide,  AS2O5,  a white,  glassy  mass.  Very  strong  ignition 
decomj)oses  this  into  arsenic  trioxide  and  oxygen  ; in  contact  with  water 
it  gradually  changes  to  arsenic  acid. 

Ortho-arsenic  acid  is  readily  soluble  in  water,  and  is  a strong  tribasic 
acid.  Its  salts — the  arseniates — are  very  similar  to  the  phosphates  and 
are  isomorphous  with  them.  With  the  soluble  salts  silver  nitrate  gives  a 
reddish-brown  precipitate  of  tri-silver  arseniate,  AgjAsO^. 


COMPOUNDS  OF  ARSENIC  WITH  SULPHUR. 

Arsenic,  like  nitrogen  and  phosphorus,  combines  in  several  proportions 
with  sulphur  : realt^ar,  AS2S2,  auripigment  or  trisulphide,  AS2S3,  and penta- 
sulphide,  AS2S5.  They  can  be  prepared  by  fusing  arsenic  and  sulphur 
together.  The  two  containing  the  more  sulphur  maybe  also  obtained  by 
passing  hydrogen  sulphide  into  solutions  of  the  arsenic  oxides.  This  indi- 
cates the  metallic  nature  of  arsenic,  for  this  is  the  universal  method 
of  preparing  metallic  sulphides: 

As.^03  -|-  3H2S  = AS2S3  + 3112^^  > AsgOg  5112^  — AS2S5  -|-  5ll2f^* 

Arsenic  Trisulphide,  AS2S3,  occurs  in  nature  as  auripigment  in  bril- 
liant yellow,  leafy,  crystalline  masses  of  specific  gravity  3.4.  It  is  pre- 
cipitated from  acidulated  solutions  of  arsenious  acid  or  its  salts  by 
hydrogen  sulphide,  as  a lemon-yellow  amorphous  powder.  It  is  insoluble 
in  water,  but  dissolves  readily  in  ammonium  hydroxide,  the  alkalies  and 
alkaline  sulphides.  It  is  worthy  of  note  that  it  dissolves  upon  prolonged 
heating  in  concentrated  hydrochloric  acid  : 

AS2S3  -|-  6IIC1  = 2ASCI3  -f  3H2S ; 

the  arsenious  chloride  then  volatilizes  with  the  hydrochloric  acid  vapors. 

Arsenic  Pentasulphide,  AS2S5,  separates  as  a bright  yellow  powder 
from  the  solution  of  sodium  sulpharseniate,  Na3AsS4  (see  below),  upon 
the  addition  of  acids.  Also  if  a rapid  current  of  hydrogen  sulphide  be 
conducted  through  a slightly  acidulated  arsenic  acid  or  arseniate  solution, 
heated  to  80°  ; the  arsenic  acid  is  then,  according  to  R.  Bunsen,  slowly 
but  completely  converted  into  ])entasuli)hide  : 

2II3ASO,  -f  51128  = AS2S3  -I  8II2O 


OXYGEN  COMPOUNDS  OF  ANTIMONY. 


223 


[Ann.  Chem.  192  (1878),  305].  Under  other  conditions  arsenic  acid 
can  be  reduced  by  hydrogen  sulphide  entirely  or  in  part  to  arseniousacid 
with  the  separation  of  sulphur : 

HaAsO^  -f  HjS  = HgAsOa  + IiP  -f  S, 

which  is  then  either  precipitated  as  trisulphide  in  the  heat  or  in  the 
presence  of  much  hydrochloric  acid  it  may  be  volatilized.  Consult 
Brauner  and  Tomicek,  Monatshefte  f.  Chem.  viii  (1887),  607;  McCay, 
Z.  f.  analyt.  Chem.  27  (1888),  632,  and  Piloty  and  Stock,  Ber.  30 
(1897),  1649,  upon  the  complex  and  as  yet  unexplained  relations. 

Arsenic  Disulphide,  As^Sg,  also  exists.  It  occurs  in  nature  as 
realgar y forming  the  beautiful,  ruby-red  crystals,  of  specific  gravity  3.5. 
It  is  applied  as  a pigment.  It  is  prepared  artificially  by  fusing  arsenic 
with  sulphur  in  the  proportion  by  weight  expressed  by  its  formula. 

Arsenic  Sulpho-salts. — Just  as  the  oxygen  compounds  of  arsenic  combine  with 
basic  oxides  and  hydroxides  to  form  salts,  so  the  sulpho-salts  are  formed  by  combination 
of  the  arsenic  sulphides  with  alkaline  sulphides  : 

'AsjSg  -j-  3K2S  = 2K,AsS3 
Tripotassium 
sulpharseuite. 

As.^Sj  -f-  3K2S  = 2K3AsS^. 

Tripotassium 

sulpharseniate. 

For  the  preparation  of  these  sulpho-salts,  arsenic  sulphide  is  dissolved  in  the  aqueous 
solution  of  the  alkaline  sulphides  or  hydrogen  sulphide  is  conducted  through  the  alkaline 
solution  of  the  oxygen  salts  : 

KgAsO^  -P  4H2S  = K3ASS4  + 4H2O. 

The  sulpho-salts  of  the  alkalies  and  ammonium  are  easily  soluble  in  water,  and  when  the 
solution  is  evaporated  they  generally  separate  in  crystals.  Acids  decompose  them,  arsenic 
sulphide  separating  out  and  hydrogen  sulphide  becoming  free  : 

2K3ASS4  -f  6HC1  ==  AS2S5  -f  6KC1  4-  3H2S. 

The  silver  salt,  Ag3AsS3,  silver  sulpharsenite,  is  the  mineral  frmistite. 

Antimony,  carbon,  tin,  gold,  platinum  and  some  other  metals  form  sulpho-salts  similar 
to  those  of  arsenic  (and  of  phosphorus). 


4.  OXYGEN  COMPOUNDS  OF  ANTIMONY. 

The  oxygen  derivatives  of  antimony  are  analogous  in  constitution  to 
those  of  arsenic:  antimony  trioxide,  Sb203,  and  antimony  pentoxide, 
81^205.  The  similarity  of  the  two  metals  is  indicated  here  just  as  in  their 
chlorides.  The  trioxide  does  not  possess  acid  but  basic  ])roperties 
almost  solely;  it  is  true  that  a few  salt-like  compounds  with  such  power- 
ful bases  as  caustic  soda  and  potash  are  known,  but  they  are  immediately 
decomposed  by  water  with  the  sejiaration  of  antimony  trioxide  oxide. 
Its  salts  with  strong  acids,  in  which  it  ai)])ears  as  a base,  are  also  easily 
decomposed  by  water.  The  normal  hydrate,  HgSbOg,  corresponding  to 


224 


INORGANIC  CHEMISTRY. 


arsenioiis  acid,  IIjAsOg,  may  be  obtained  from  ])otassinm  antimonyl  tar- 
trate (see  Organic  Chemistry).  Other  hydrates  are  not  definitely  known. 

The  higher  oxidation  product,  the  i)entoxide,  Sl)/)^,  on  the  contrary, 
has  an  acid  nature  and  yields  salts  with  the  bases.  The  hydrate,  IljSbtJ^, 
or  ortho-antimonic  acid,  ])asses  at  175°  into  vieta-antimonic  acid,  HSbOg, 
and  at  275°  into  the  anhydride,  Sb./)^.  Most  of  the  antimoniates  are 
derived  from  pyro-antimonic  acid,  H^Sb^O,,  which  is  produced  on  decom- 
posing antimony  pentachloride  with  water  and  drying  the  precii)itate  at 
100°  (p.  148). 

Antimony  Oxide  (antimony  trioxide),  SbjOj  or  Sb^Og,  is  obtained  by  burning  the 
metal  in  the  air,  or  by  oxidizing  it  with  dilute  nitric  acid.  By  sublimation  it  may  be 
obtained  in  two  different  crystal  systems,  in  regular  octahedra  and  in  rhombic  prisms.  In 
these  forms  it  occurs  also  in  nature  as  the  isometric  senarniontite  and  rhombic  valentinite. 
It  is  not  isodimorphous  with  arsenic  trioxide,  which  was  long  considered  to  be  the  case. 
Its  vapor  density  at  1560°  corresponds  to  the  formula  Sb^Og,  and  at  higher  temperatures 
it  is  very  probable  that  it  will  be  like  arsenic  trioxide,  the  simpler  molecule  Sb203.  d'he 
hydrate,  Sb(OII)3,  is  thrown  out  as  a white  precipitate  on  adding  dilute  sulphuric  acid 
to  a solution  of  tartar  emetic.  It  parts  with  water  quite  readily  and  changes  to  antimony 
trioxide  also  called  antimonious  acid. 

The  latter  and  the  hydrate  dissolve  in  sodium  and  potassium  hydroxide,  and,  very 
probably,  form  salts  (NaSb02)  which  decompose  upon  evaporating  the  solution.  In  this 
beliQvior  the  acid  nature  of  antimony  hydrate  is  also  seen  ; therefore  it  has  received  the 
name  of  antimonious  acid. 

The  oxide  forms  salts  with  acids,  which  are  derived  either  from  the  normal  hydrate, 
HjSbO^,  or  from  the  hydrate,  HSbOj  = SbO.OII  {ineia-antinionions  acid).  In  the 
salts  of  the  first  kind  we  have  three  hydrogen  atoms  of  the  hydrate  replaced  by  acid 
radicals,  or,  what  is  the  same,  a trivalent  antimony  atom  displacing  three  hydrogen  atoms 
of  the  acids : 

Sb03fN02)3  or  Sb(N03)3. 

Antimony  nitrate. 

In  the  second  variety  of  antimony  salts  derived  from  the  hydrate,  SbO . OH,  hydrogen 
is  replaced  by  a univalent  acid  residue,  or  the  hydrogen  of  the  acid  is  substituted  by  the 
univalent  group,  SbO,  known  as  antimonyl  : 

SbO.O.NO^  or  N03(Sb0). 

Antimonyl  nitrate. 

Of  these  salts  may  be  mentioned  the  following  : 

Antimony  Sulphate,  Sb2(S04)3,  which  separates  when  a solution  of  the  oxide  in 
concentrated  sulphuric  acid  is  cooled. 

Antimonyl  Sulphate,  (Sb0)2S04,  is  formed  when  antimony  oxide  is  dissolved  in 
somewhat  dilute  .sulphuric  acid,  and  crystallizes  in  fine  needles  on  cooling.  Water  decom- 
poses both,  formii>g  basic  salts  ; hence  the  basic  nature  of  antimony  oxide  is  slight. 

Antimonic  Acid,  H3SbO^,  is  obtained  upon  warming  antimony  with  concentrated 
nitric  acid  or  by  adding  antimcrny  pentachloride  to  cold  water  : 

2SbCl5  -f  8II2O  = 2ll3SbO,  + loHCl. 

It  is  most  conveniently  prej)ared  by  decomposing  its  potassium  salt  with  nitric  acid.  It 
forms,  when  dried  over  concentrated  .sulphuric  acid,  a white  powder  of  the  composition 
indicated  by  the  formula.  When  air-dried  it  has  the  formula — 

2n3SbO^  T II2O. 

It  is  almost  insoluble  in  water  and  in  nitric  acid,  but  reddens  blue  litmus-paper.  It  is  a 
weak  monobasic  acid,  the  .salts  oi'  which  are  mostly  insoluble  in  water.  [See  Beilstein 


COMPOUNDS  OF  ANTIMONY  WITH  SULPHUR. 


225 
See  also 


and  V.  Blaese,  Cheni.  Centralblatt,  1889,  i,  803;  Ebel,  Ber.  22  (1889),  3044. 
p.  225.] 

At  4(X)°  antimonic  acid  yields  antimony  pentoxide,  Sb.^Oj,  a yellow,  amorphous 
mass,  soluble  in  hydrochloric  acid.  At  800°  the  pentoxide  breaks  down  into  oxygen  and 
the  oxide,  Sb204,  which  can  be  viewed  as  antimonyl  meta-antimoniate  (SbOj.  SbO),  or 

as  a mixed  anhydride,  | O.  It  is  a white  powder,  becoming  yellow  when  heated 
and  at  the  melting  point  of  silver  is  decomposed  into  oxygen  and  the  trioxide. 


COMPOUNDS  OF  ANTIMONY  WITH  SULPHUR. 

These  are  perfectly  analogous  to  the  sulphur  compounds  of  arsenic,  and 
form  sulpho  salts  with  alkaline  sulphides,  corresponding  to  the  oxygen 
salts.  Acids  precipitate  antimony  sulphide  and  liberate  hydrogen  sul- 
phide from  the  sulpho-salts. 

Antimony  Trisulphide,  Sb2S3,  is  found  in  nature  as  stibnite  {Stibium 
sulphuratu77i  nigriwi)  in  radiating  crystalline  masses  of  dark-gray  color 
and  metallic  luster;  its  specific  gravity  is  4.7.  When  heated  it  melts  and 
distils.  The  artificial  sulphide  obtained  by  precipitating  a solution  of 
the  oxide  with  hydrogen  sulphide,  is  an  amorphous  red  powder.  When 
it  is  digested  with  aqueous  hydrochloric  acid  or  when  carefully  heated 
alone  in  an  atmosphere  of  carbon  dioxide  it  passes  into  the  black,  crys- 
talline variety.  Upon  warming,  the  sulphide  dissolves  in  concentrated 
hydrochloric  acid  to  form  antimony  trichloride. 

The  compound,  Sb2S20,  occurring  in  nature  as  red  stibnite,  can  be 
artificially  prepared,  and  serves  as  a beautiful  red  pigment,  under  the 
name  of  a77timony  ciTiTiabar.  Ke7  77ies  77imerale,  employed  in  medicine,  is 
obtained  by  boiling  antimony  sulphide  with  a sodium  carbonate  solution, 
and  is  a mixture  of  antimony  trisulphide  and  antimony  trioxide. 

Antimony  Pentasulphide,  S^S^,  or  gold  sulphur  (Stibiu77t  sulphur- 
atu77i  aurantiacuni),  is  precipitated  by  hydrogen  sulphide  from  acid  solu- 
tions of  antimonic  acid  ; it  is  more  conveniently  obtained  by  the  precipi- 
tation of  sodium  sulphantimoniate,  Na^SbS^,  with  hydrochloric  acid: 

2Na3SbS,  -f  6IIC1  = Sb2S5  + 6NaCl  + 3H2S. 

It  is  an  orange-red  powder,  like  the  trisulphide;  it  decomposes  on 
being  heated  into  antimony  trisulphide  and  sulphur.  It  dissolves  to 
antimony  trichloride  in  strong  hydrochloric  acid,  with  separation  of  sul- 
phur and  liberation  of  hydrogen  sulphide. 

Sodiu77i  sulpha7tti77io7iiatey  NajSbS^  (Schlippe’s  salt),  results  from  boiling 
pulverized  antimony  trisulphide  with  sulphur  and  sodium  hydroxide 
(p.  223).  Upon  concentrating  thesolution  it  crystallizes  in  large  yellow 
tetrahedra  containing  nine  molecules  of  water  (SbS^Naj  -(-  pHjO) ; ex- 
posed to  the  air  it  becomes  covered  with  a brown  layer  of  antimony 
pentasulphide.  It  serves  principally  for  the  preparation  of  the  officinal 
gold  sulphur. 


226 


INORGANIC  CHEMISTRY. 


VANADIUM.  NIOBIUM.  TANTALUM. 

V = 5i.2.  Nb  = 94.  Ta  = 183. 

The  three  rare  elements,  vanadium,  niobium  and  tantalum,  are  closely  related  to  the 
phosphorus  grou^).  They  yield  derivatives  very  much  like  those  of  the  phosphorus 
group,  but  possess  a more  metallic  character.  'I'hey  exhibit  many  characteristics  similar 
to  those  of  chromium,  iron  and  tungsten,  with  which  they  are  fref[uently  associated  in 
their  naturally  occurring  compounds  (compare  the  Periodic  System  of  the  Elements). 

Vanadium,  observed  in  1801  by  Del  Rio  and  considered  to  be  chromium  until  Sefstrotn 
in  1830  proved  it  to  be  a new  element,  occurs  in  nature  principally  in  the  form  of  salts 
of  vanadic  acid  (vanadium  lead  ore).  Creuzot  has  recently  worked  the  'I'homas  slags  for 
vanadic  acid,  which  originally  was  present  in  the  iron  ores.  Vanadium  may  be  obtained 
free  by  igniting  its  chlorides  in  a current  of  hydrogen.  It  is  a grayish- white,  metallic, 
lustrous  powder,  of  specific  gravity  5.5.  It  is  difficultly  fusible,  and  changes  in  the  air 
very  slowly  at  ordinary  temperature.  When  heated,  it  burns  to  vanadium  pentoxide, 
VjOj.  It  unites  readily  with  nitrogen  to  form  vanadium  mononitride,  VN. 

Vanadium  trichloride,  VCI3,  forms  red  plates,  which  readily  deliquesce  in  the  air; 
it  is  not  volatile. 

Vanadium  oxychloride,  VOCI3,  results  on  heating  a mixture  of  vanadium  trioxide, 
V,^03,  and  carbon  in  chlorine  gas.  It  is  a lemon-yellow  liquid,  of  specific  gravity  1.84, 
and  boils  at  126°.  It  fumes  strongly  in  the  air  and  decomposes  with  water  (analogous 
to  phosphorus  oxychloride)  into  vanadic  acid  and  hydrochloric  acid.  Its  vapor  density 
corresponds  to  the  formula  VOCI3. 

Vanadium  trioxide,  V.2O3,  is  a black  powder  obtained  by  heating  vanadium  pentoxide, 
V2O5,  in  hydrogen.  It  combines  with  oxygen,  to  form  vanadium  pentoxide,  VgO^.  The 
corresponding  sulphate,  V2(SO^)3,  combines,  like  the  similarly  constituted  sulphates  of 
aluminium,  iron  and  chromium,  with  the  sulphates  of  the  alkalies  to  form  alums  ( see  these ). 

Vanadium  pentoxide,  V or  vanadic  anhydride,  is  a brown  mass  obtained  by  fusing 
the  naturally  occurring  vanadates  with  niter,  etc.  It  exists  in  three  varieties.  It  is  solu- 
ble in  the  alkalies,  and  forms  with  the  metals  salts  of  vaiiadic,  Il3V()4,  and  metavanadic 
acids,  HVO3.  All  these  compounds  are  similar  in  constitution  to  those  of  the  elements 
of  the  phosphorus  group.  In  addition  to  these,  vanadium  forms  other  compounds,  con- 
stituted like  those  of  sulphur  and  chromium.  In  this  class  belong  VCI2  (dichloride), 
the  tetrachloride,  VCl^,  vanadious  oxide,  VO,  vanadium  dioxide,  VO2,  and  the  oxychlo- 
ride, VOClg.  The  tetrachloride,  V.C1^,  is  a red-brown  liquid,  boiling  at  154°  ; its  vapor 
density  corresponds  to  its  formula. 

Niobium,  Nb,  and  tantalum,  Ta,  are  not  well  known  in  the  free  state.  They  occur 
together  as  niobates  and  tantalates  in  a few  rare  minerals — the  columbites  and  tantalites. 
The  chlorides,  NbClg  and  TaCl^,  are  volatile  and  are  decomposed  by  water.  Niobium 
and  tantalum  unite  with  potassium  fluoride,  forming  double  salts,  e.g.,  2KFl.NbFl5  and 
2KFl.TaFl5;  also  2KFI . NbOFlg  and  2KFI . TaOFlg.  When  potassium  niobium  fluor- 
ide, 2KFI.  NTFI5,  is  heated  with  sodium,  niobium  hydride,  NbH,  is  formed.  This  is  a 
grayish-black  powder,  which  when  heated  burns  to  niobic  anhydride,  Nb205,  and  water. 
The  oxides,  Nb205  and  form  salts  of  niobic  (HgNbO^)  and  tantalic  (HgTaO^) 

acids  with  bases. 


4.  OXYGEN  DERIVATIVES  OF  THE  ELEMENTS  OF 
THE  CARBON  GROUP. 


'File  following  normal  hydroxides  correspond  to  the  halogen  deriva- 
tives, CCl^,  SiCl^,  GeCl^,  and  SnCl^,  of  the  quadrivalent  elements  car- 
bon, silicon,  germanium,  and  tin  (see  p.  164): 


IV  IV 

C(()II)^  Si(OlI)^ 

Normal  Normal 

carbonic  acid.  silicic  acid. 


TV 

Oe(OII), 

Normal 

germanic  acid. 


IV 

Sn(OH)4. 

Normal 
stannic  acid. 


OXYGEN  COMPOUNDS  OF  CARBON. 


227 


These  normal  hydrates  or  acids  being  true  ortho-acids  have  but  little 
stability,  and  exist  chiefly  in  their  derivatives.  By  the  separation  of  one 
molecule  of  water,  they  pass  into — 

CO3H2  SiOgH^  GeOgHj  SnOgllg, 

Or 

C0(0H)2  SiO(OH),  GeO(OH)2  SnO(OH)2. 

Carbonic  acid.  Silicic  acid.  Germanic  acid.  Stannic  acid. 

These  hydroxyl  derivatives  deport  themselves  toward  the  normal  just 
as  the  meta  acids  of  the  elements  of  the  nitrogen  group  do  to  the  ortho- 
acids (see  p.  201).  They  constitute  the  ordinary  acids  of  the  quadriv- 
alent elements  carbon,  silicon,  germanium,  and  tin,  and,  as  they  contain 
two  hydroxyl  groujDS,  are  dibasic. 

Carbon  is  the  lowest  member  of  this  group,  with  the  least  atomic 
weight.  Among  the  elements  of  the  other  three  groups  corresponding  to 
it  are  nitrogen,  oxygen,  and  fluorine : 

C = 12  N = 14.04  O = 16  FI  = 19. 

Fluorine  and  oxygen  do  not  afford  any  oxygen  acids.  The  normal  acids 
of  nitrogen,  N(OH)5  and  N(OH)3,  are  very  unstable,  and  pass  readily 
into  the  meta-acids,  NO2 . OH  and  NO . OH.  The  normal  carbonic  acid, 
C(OH)^,  corresponds  to  this,  but  is  not  capable  of  existing.  Indeed,  the 
meta-  or  ordinary  carbonic  acid,  H2CO3,  is  also  very  unstable  and  at  once 
decomposes,  when  separated  from  its  salts,  into  water  and  carbon  dioxide, 
CO2.  Even  silicic,  germanic,  and  stannic  acids  break  down  readily  into 
water  and  their  anhydrides: 

COg  SiOg  GeOg  SnOg. 

Carbon  dioxide.  Silicon  dioxide.  Germanium  dioxide.  Tin  dioxide. 


1.  OXYGEN  COMPOUNDS  OF  CARBON. 

Carbon  Dioxide,  CO2,  or  carbonic  anhydride  (generally  called  car- 
bonic acid),  is  produced  when  carbon  or  its  compounds  are  burned  in  air 
or  oxygen.  It  is  found  free  in  the  air  (in  100  volumes,  on  an  average, 
0.035  volumes  CO2),  in  many  mineral  springs  (acid  springs),  and  escapes 
in  large  quantities  from  the  earth  in  many  volcanic  districts.  It  occurs 
in  the  liquid  form,  enclosed  in  the  cavities  of  many  crystalline  minerals 
(quartz,  topaz).  It  is  prepared  on  a large  scale  by  burning  coke  or  by 
the  ignition  of  limestone;  in  the  laboratory  it  may  be  most  conveniently 
obtained  by  the  decomposition  of  calcium  carbonate  (marble  or  chalk) 
with  dilute  hydrochloric  acid  : 

CaCOg  + 2IICI  = CaCl.2  + COg  + ITgO. 

Calcium  Calcium 

carbonate.  chloride. 

Carbon  dioxide  is  a colorless  gas,  with  a slightly  acid  taste.  Owing 
to  its  weight,  the  gas  may  be  collected  by  air  disj)lacement,  and  may  be 


228 


INORGANIC  CHEMISTRY. 


l)oured  from  one  vessel  into  another  filled  vvitli  air.  A liter  of  carbon 
dioxide  weighs  i 965  grams  at  0°,  760  mm.  pressure,  and  in  45°  latitude 
at  sea-level.  By  a pressure  of  50-60  atmospheres  at  the  ordinary  tem- 
jierature  carbon  dioxide  can  be  licpiefied,  as  was  first  shown  by  h'araday. 
'I'he  apparatus  of  Thilorier  and  Natterer  were  emjjloyed  to  this  end.  At 
l)resent  licpiid  carbon  dioxide  is  brought  into  market  enclosed  in  wrought- 
iron  cylinders,  and  is  used  (piite  regularly  in  technical  operations. 

Carbon  dioxide  can  only  be  licpicfied  below  -1-30.9°;  this  is  its  criti- 
cal temperature  (p.  47).  Its  tension  (critical  pressure)  at  this  point 
ecpials  77  atmosi)heres.  If  licpiid  carbon  dioxide,  enclosed  in  some  suit- 
able vessel,  be  allowed  to  escajie  into  the  air  by  opening  a stop-cock 
(ordinary  pressure),  it  immediately  solidifies  (see  below)  to  a white,  snowy 
mass.  This  is  because  in  the  eva[)oration  of  a jiart  of  the  licpiid  .so  much 
heat  is  withdrawn  that  the  remainder  becomes  solid.  vSolid  carbon 
dioxide  is  a very  poor  conductor  of  heat  and  vajiorizes  very  slowly.  Not- 
withstanding its  low  tem])erature  it  can  be  handled  without  serious  result, 
because  it  is  always  surrounded  by  a gaseous  layer,  and  is,  consecpiently, 
not  in  immediate  contact  with  the  skin.  If,  however,  it  be  jiressed 
between  the  fingers,  it  will  produce  painful  burns. 

The  temperature  of  the  solid  carbon  dioxide  vaporizing  in  the  air  under 
ordinary  pressure  is  about  — 78°  (its  boiling  point).  When  the  solid 
dioxide  is  mixed  with  a little  ether  it  forms  a paste,  and  then  conducts 
heat  better,  and  is,  therefore,  well  adapted  as  a cooling  agent.  In  vacuo 
its  temperature  diminishes  to  — 140°. 

Liquid  carbon  dioxide  is  a colorless,  very  mobile  liquid.  Its  specific 
gravity  is  0.91  at  — 1.6°,  0.84  at  -|-i5°,  and  0.726  at  -[-22.2°.  Its  coeffi- 
cient of  expansion  is,  consequently,  greater  than  that  of  the  gases  ; other 
gases  behave  similarly,  but  only  such  as  are  condensed  under  great  pressure. 
If  liquid  carbon  dioxide,  contained  in  a glass  tube,  be  heated,  it  expands 
rapicily  and  suddenly  passes,  at  the  critical  temperature,  -{-30-9°)  ii^to 
gas.  This  behavior  enables  us  to  determine  without  difficulty  whether 
the  liquids  contained  (see  above)  in  minerals  are  liquid  carbon  dioxide. 

If  liquid  carbon  dioxide,  confined  in  a glass  tube,  be  cooled  by  a mix- 
ture of  solid,  snowy  dioxide  and  ether  (see  above),  it  will  solidify  to  a 
transparent  ice-like  mass,  which  will  melt  at  — 65°,  according  to  another 
authority  at  — 57°. 


The  tension  of  the  solid  or  liquid  dioxide,  which  at  the  same  time  indicates  the  pres- 
sure necessary  for  condensation  at  various  temperatures,  is  given  in  the  following  table  : 


Temperaturk. 

Tension. 

Temperature. 

Tension. 

4 30.9° 

77  atmos. 

— 20° 

19.9  atmos. 

20° 

58.0  “ 

0 

0 

1 

10.2  “ 

10° 

46.0  “ 

—60° 

3-9  “ 

0° 

35-4  “ 

— 70° 

2.1  “ 

—80° 

I.O  “ 

CARBON  DIOXIDE. 


229 


At  the  temperature  of  fusion  of  the  solid  dioxide  ( — 65°)  the  tension  equals  about  3.5 
atmospheres  ; the  resulting  liquid  has  this  tension  at  this  temperature.  If  the  external 
pressure  exerted  upon  it  be  less,  it  cannot  exist  as  a liquid,  but  must  immediately  pass 
into  the  gaseous  state.  Herein  we  observe  why  the  solid  dioxide  (under  ordinary  pres- 
sure) does  not  melt  in  the  air,  but  vaporizes  at  once  ; and  further,  it  explains  why  the 
liquid  dioxide,  subjected  to  the  ordinary  atmospheric  pressure,  cannot  continue  in  this 
state — why  it  is  either  gasified  at  once  or  changed  to  the  solid  form. 

Many  other  fusible  solids  behave  like  the  dioxide.  If  the  tension  of  their  vapors  at 
the  fusion  temperature  exceeds  that  of  the  external  atmospheric  pressure  they  do  not  melt 
in  the  air,  because  the  resulting  liquid  is  immediately  transformed  into  vapor ; they 
vaporize  (sublime)  directly,  without  previous  fusion.  Such  bodies  are,  e.  g.,  arsenic, 
arsenic  trioxide,  AsjOg,  camphor,  mercurous  chloride,  HgCl,  etc.  They  can  be  fused  only 
under  increased  pressure  (in  sealed  tubes).  Again,  all  solids  fusible  in  the  air  (under 
ordinary  pressure)  may  be  converted  directly  into  gases  by  removing  the  external  pres- 
sure. Thus  iodine  fuses  at  114°,  but  sublimes  in  a vacuum  without  previous  fusion. 
Mercuric  chloride,  HgClj,  fuses  at  265°,  but  not  if  the  external  pressure  be  less  than  420 
mm.  Water  melts  at  0°,  its  tension  at  this  temperature  is  4. 6 mm.  If  the  external  pres- 
sure be  less  {in  vacuo),  it  will  no  longer  melt,  but  vaporize  at  once.  The  pressure 
below  which  solids  no  longer  melt  has  been  called  Xhexx  critical  pressure  (Carnelley).  It 
is,  of  course,  understood  that  this  is  nothing  more  than  the  tension  of  the  substance  at  its 
point  of  fusion. 

Water  dissolves  an  equal  volume  of  the  gas  at  1 4°  ; at  0°  it  takes  up  i . 79 
volumes.  This  proportion  remains  constant  for  every  pressure,  i.  e.,  at 
every  pressure  the  same  volume  of  gas  is  absorbed.  As  gases  are  con- 
densed in  proportion  to  the  pressure,  the  quantity  of  absorbed  gas  is  also 
proportional  to  the  former  (law  of  Henry  and  Dalton).  Hence  i volume  of 
water  absorbs,  at  14°  and  2 atmospheres  pressure,  2 volumes,  at  3 atmos- 
pheres 3 volumes,  etc.,  of  carbon  dioxide — measured  at  ordinary  pressure. 
The  gas  absorbed  at  higher  pressure  escapes  with  effervescence  of  the  liquid 
when  the  pressure  is  diminished;  upon  this  depends  the  sparkling  of  soda 
water  and  champagne,  which  are  saturated  with  carbon  dioxide  under 
high  pressure.  Every  naturally  occurring  water,  especially  spring  water, 
holds  carbon  dioxide  in  solution,  which  imparts  to  it  a refreshing  taste. 

As  the  product  of  a complete  combustion  carbon  dioxide  is  not  com- 
bustible, and  is  unable  to  support  the  combustion  of  most  bodies;  a glim- 
mering chip  is  immediately  extinguished  in  it.  In  a similar  manner  it  is 
irrespirable.  Although  it  is  not  poisonous,  in  the  true  sense  of  the 
word,  yet  the  admixture  of  a few  per  cent,  of  carbon  dioxide  to  the  air 
makes  it  suffocating,  as  it  retards  the  elimination  of  the  same  gas  from  the 
lungs. 

It  is  decomposed  by  the  continued  action  of  the  electric  spark  into  car- 
bon monoxide  (CO)  and  oxygen ; upon  heating  to  1300°  it  suffers  a par- 
tial decomposition  into  carbon  monoxide  and  oxygen.  It  is  also  decom- 
posed when  conducted  over  heated  potassium  or  sodium,  with  separation 
of  carbon  ; the  potassium  combines  with  oxygen  to  form  potassium  oxide  : 


CO2  -f-  2K2  C -j-  2K2O, 

which  forms  potassium  carbonate  (K.^C03),  with  the  excess  of  carbon  di- 
oxide. Glowing  carbon  reduces  the  dioxide  to  the  monoxide  (p.  231). 
Carbon  monoxide  is  analogously  formed  on  conducting  a mixture  of  the  di- 
oxide and  hydrogen  (equal  volumes)  through  a tube  heated  to  redness,  but 


230 


INORGANIC  CHEMISTRY. 


between  250°  and  300°  the  opposite  reaction  takes  place,  although  only 
to  a slight  degree  : 

CO,  + n,  ^ CO  + 11,0. 

The  composition  of  carbon  dioxide  is  readily  determined  by  burning 
a weighed  quantity  of  pure  carbon  (diamond  or  graphite)  in  a current 
of  oxygen,  and  ascertaining  the  weight  of  the  resulting  gas.  From  the 
formula  COj  it  follows  that  in  one  volume  of  carbon  dioxide  there  is 
contained  an  equal  volume  of  oxygen.  We  may  satisfy 
ourselves  of  this  by  burning  carbon  in  a definite  volume 
of  oxygen;  after  cooling,  there  is  obtained  an  equal 
volume  of  carbon  dioxide  : 

C + O2  = COj. 

I vol.  I vol. 

The  experiment  is  most  i)ractically  executed  by  aid  of 
the  apparatus  of  Hofmann  pictured  in  Fig.  66.  The 
sphere-sha})ed  expansion  of  the  eudiometer  limb  of  the 
U-tube  is  closed  by  means  of  a glass  stopper,  through  which 
two  copper  wires  pass.  The  one  wire  bears  a combustion 
spoon  at  its  end,  upon  which  lies  the  carbon  to  be  burned, 
while  the  other  terminates  in  a thin  piece  of  platinum, 
which  is  in  contact  with  the  carbon.  For  the  performance 
of  the  experiment,  the  [J-tube  is  filled,  to  near  the  bulb, 
with  mercury  and  the  air  is  expelled  from  the  bulb  limb 
by  means  of  a rapid  current  of  oxygen,  the  stapj)er  made 
air-tight,  the  mercury  level  noted,  and  then  pass  the  in- 
duction spark  between  the  platinum  wire  and  the  copper 
Fig.  66.  spoon;  this  induces  the  burning  of  the  carbon.  As  the 
volume  of  the  enclosed  gas  is  greatly  expanded  by  the  heat 
developed,  it  is  advisable,  in  order  to  avoid  the  jumping  out  of  the  stopper, 
to  previously  reduce  the  pressure  of  the  gas  about  two-thirds,  by  running 
out  mercury.  When  the  combustion  is  finished  the  mercury  which  was 
drawn  off  is  returned  and  it  rapidly  assumes  its  original  position.  The 
same  apparatus  can  also  be  employed  for  the  illustration  of  the  volume 
relations  observed  in  the  combustion  of  sulphur  and  other  bodies. 

In  dry  condition,  carbon  dioxide,  like  all  anhydrides,  exhibits  neither 
basic  nor  acid  reaction.  In  aqueous  solution  it  colors  blue  litmus-paper 
a faint  red  ; upon  drying  the  paper  the  red  disappears,  in  consequence 
of  the  evaporation  of  the  carbon  dioxide. 

We  may  then  regard  it  as  probable  that  free  carbonic  acid,  H^COj,  is 
contained  in  the  aqueous  solution,  but  this  readily  decomposes  into  the 
dioxide,  COj,  and  water.  The  salts  of  carbonic  acid  are  produced  by 
the  action  of  carbon  dioxide  upon  the  bases : 

2KOII  + CO2  = K2CO3  -f  HjO. 

Potassium 

carbonate. 

Carbon  dioxide  is,  therefore,  easily  absorbed  by  potassium  and  sodium 
hydroxide.  On  conducting  it  through  a solution  of  calcium  or  barium 


CARBON  MONOXIDE.  23 1 

hydroxide,  a white  precipitate  of  barium  or  calcium  carbonate,  BaCO^ 
or  CaCOg,  is  produced. 

Carbonic  acid  is  dibasic,  forming  primary  (acid)  and  secondary  (neu- 
tral) salts,  KHCO3  and  K2CO3,  called  carbonates.  As  the  acidity  of  car- 
bonic acid  is  only  slight,  the  secondary  salts,  obtained  from  strong  bases, 
exhibit  a basic  reaction.  Most  acids  expel  the  weak  carbonic  acid  from 
its  salts,  with  decomposition  into  carbon  dioxide  and  water. 

In  the  reduction  of  carbon  dioxide  we  meet  an  interesting  transi- 
tion or  transformation  of  substances  of  mineral  origin  into  those  of  ani- 
mate nature.  According  to  Lieben  nascent  hydrogen  readily  reduces 
carbon  dioxide  in  the  form  of  dissolved  bicarbonates  to  formic  acid  (for- 
mates) : 

HO  . COONa  + = H . COONa  + H^O. 

Sodium  Sodium 

bicarbonate.  formate. 

And  the  latest  researches  show  that  water  and  hydrogen,  under  the  influ- 
ence of  the  silent  electric  discharge,  reduce  carbon  dioxide  to  formic 
acid  : 

CO.2  + H2  = H . COOH  and  CO2  + H2O  = H . COOH  + O. 

The  reduction  of  carbon  dioxide  by  the  chlorophyl  granules  of  plants 
in  the  sunlight  is  ])articularly  important.  It  is  by  them  that  the  almost 
endless  number  of organic  bodies  ” contained  in  plants  is  built  up.  In 
the  animal  organism,  on  the  other  hand,  carbon  dioxide  and  water  are 
the  chief  decomposition  products  of  organic  compounds ; consequently 
the  exhaled  air  is  rich  in  this  acid. 

Percarbonic  acid,  H2C.20g,  has  been  obtained  by  the  electrolysis  of  acid  carbonates  just 
as  persuiphuric  acid  was  formed  from  sulphuric  acid.  It  is  due  to  the  union  of  the  ions — 

O-CO-OH 

O-CO-Me  (Me  = metal).  The  free  acid  would  have  the  formula  ) 

O-CO-OH’ 

and  may  be  said  to  be  derived  from  hydrogen  peroxide  (p.  198  ).  Its  salts  will  be 
described  later. 

Carbon  Monoxide,  CO.  Carbon  dioxide  is  the  first  product  in  the 
combustion  of  coal.  But  just  as  soon  as  the  combustion  reaches  a certain 
temperature — maximum — the  excess  of  carbon  reduces  the  dioxide  and 
carbon  monoxide  is  formed  : 

CO2  -f  C = 2CO. 

1 vol.  2 vols. 

This  is  demonstrated  by  the  following  experiment : When  dry  air  is  con- 
ducted over  heated  coal,  carbon  dioxide  is  formed  almost  exclusively 
from  400°  to  about  700°.  P'rom  this  point  forward  carbon  monoxide 
appears  in  increasing  amount  and  from  1000°  upwards  it  is  almost  the 
sole  product.  Hence  glowing  coals  at  moderate  temperature  burn  with- 
out flame,  but  at  more  elevated  temperatures  (from  1000°  upwards)  flame 
is  present  (p.  155  ). 

Zinc-dust  reacts  like  carbon  : 


COj  + Zn  = ZnO  -f  CO. 


232 


INORGANIC  CHEMISTRY. 


When  carbon  dioxide  is  conducted  through  a glass  tube,  containing  zinc-dust  heated  to  a 
faint  red  heat,  almost  pure  carbon  monoxide  escapes.  A more  convenient  procedure 
consists  in  heating  pulverized  magnesium  carbonate  and  zinc-dust  in  a gla.ss  retort,  when 
CO,  containing  COj,  is  eliminated  ; .subsequently  the  former  alone  e.scapes.  Pure  mon- 
oxide is  al.so  formed  upon  heating  zinc-dust  with  chalk  (in  Cfjui valent  quantities)  : 

Zn  + CaCOj  = ZnO  -f  CaO  -j-  CO, 
a very  convenient  method  for  its  preparation. 

The  monoxide  is  prodticed,  further,  by  igniting  carbon  with  different 
metallic  oxides,  e.  g.,  zinc  oxide : 

ZnO  -P  c = Zn  -f  CO. 

Water  is  similarly  decomposed.  On  conducting  aqueous  vapor  over 
burning  carbon  there  is  produced  a mixture  of  carbon  dioxide  and 
hydrogen : 

C -f  2lbO  = COj  + 2ll.,. 

By  further  action  of  the  heated  carbon  the  carbon  dioxide  in  this  equa- 
tion is  reduced  to  carbon  monoxide.  This  gas  mixture  is  known  as  water 
gas  and  is  applied  technically.  [See  A.  Naumann,  Ber.  25  (1892),  i, 
556;  also  p.  229.] 

A mixture  of  hydrogen  and  carbon  monoxide  is  also  produced  when 
the  electric  arc  passes  between  carbon  points  under  water : 

HjO  + C = CO  -f  Hj. 

For  the  preparation  of  carbon  monoxide,  oxalic  acid  is  warmed  with 
sulphuric  acid  ; the  latter  withdraws  water  from  the  former,  and  the  residue 
breaks  down  into  carbon  dioxide  and  monoxide  : 

H2C2O,  = CO2  + CO  + H2O. 

The  disengaged  mixture  of  gases  is  conducted  through  an  aqueous 
solution  of  sodium  hydroxide,  by  which  the  carbon  dioxide  is  absorbed, 
the  monoxide  passing  through  unaltered.  Pure  carbon  monoxide  may  be 
prepared  by  heating  yellow  prussiate  of  potassium  (see  Iron)  with  9 parts 
of  concentrated  sulphuric  acid.  The  resulting  gas  is  conducted  through 
sodium  hydroxide  to  remove  from  it  traces  of  carbon  dioxide  and  sulphur 
dioxide.  Pure  monoxide  is  also  produced  when  concentrated  formic  acid 
or  lead  formate  is  heated  with  concentrated  sulphuric  acid: 

CO2H2  = CO  + H2O. 

A liter  of  the  gas  weighs  1.25078  grams  under  normal  conditions;  it 
is  therefore  0.9672  as  heavy  as  air.  It  is  one  of  the  gases  which  are 
condensed  with  difficulty.  Its  critical  temperature  is  — 141°  and  its 
critical  pressure  is  35  atmospheres.  Liquid  carbon  monoxide  solidifies 
under  100  mm.  ])ressure  at  — 207°  and  under  4 mm.  pressure  at  — 220° 
(Olszewsky).  It  boils  at  — 190°  under  760  mm.  j)ressure.  It  is  almost 
insoluble  in  water,  but  is  readily  dissolved  by  an  ammoniacal  solution  of 
cui)rous  chloride  (CU2CI2)  with  which  it  forms  a crystalline  compound, 
but  this  decomposes  when  its  solution  is  heated  and  carbon  monoxide  is 


CARBON  MONOXIDE. 


233 


again  liberated.  When  ignited,  it  burns  in  the  air,  with  a faintly  lumi- 
nous, beautiful  blue  flame,  which  distinguishes  it  from  other  combustible 
gases.  With  air  or  oxygen,  it  affords  (similar  to  hydrogen)  a very  ex- 
plosive mixture  : 

2CO  + 02  = 2CO2. 

2 vols.  I vol.  2 vols. 

The  union  of  carbon  monoxide  and  oxygen  takes  place  at  very  high 
temperatures;  hence  the  burning  flame  of  the  gas  is  extinguished  upon 
cooling.  A flame  or  a spark  from  a powerful  induction  coil  is  necessary 
to  ignite  a dry  mixture  of  carbon  monoxide  and  oxygen.  When  the  two 
gases  are  moist  they  are  more  easily  ignited  and  combustible.  This  is 
explained  by  the  fact  that  carbon  monoxide  unites  with  the  aqueous 
vapor  and  yields  the  dioxide  and  hydrogen  (CO  -|-  H2O  = CO2  -j-  H2), 
which  then  combines  with  oxygen  and  forms  water  (pp.  loi,  230). 

In  consequence  of  its  ready  oxidation,  it  is  capable  of  reducing  most 
metallic  oxides  at  a red  heat : 

CuO  -4-  CO  = Cu  -h  CO2  . . . + (31.3  Cal.) 

(37.1  Cal.)  (28.5  Cal.)  (96.9  Cal.) 

Some  noble  metals  are  precipitated  from  solutions  of  their  salts  by  car- 
bon monoxide  even  in  the  cold.  Thus,  palladium  and  gold  are  thrown 
out  from  their  chloride  solutions  by  it.  A piece  of  paper  moistened  with 
palladious  chloride  (PdClg)  is  blackened  by  it  (delicate  test  for  CO). 
When  carbon  monoxide  is  conducted  into  a platinic  chloride  solution 
platinous  chloride  separates. 

Carbon  monoxide  is  reduced  to  carbon  with  difficulty.  Burning  bodies 
are  extinguished  by  it.  When  heated  with  potassium  it  is  decomposed 
with  separation  of  carbon.  Under  the  influence  of  the  silent  electric 
discharge  hydrogen  reduces  it  to  formic  aldehyde: 

CO  + H2  = H . COH. 

This  is  the  simplest  organic  compound,  consisting  of  carbon,  oxygen  and 
hydrogen.  Water  under  like  conditions  changes  carbon  monoxide  to 
formic  acid : 

CO  + II2O  = H.COOH 

(see  p.  231). 

When  inhaled,  even  in  slight  quantity,  it  acts  very  poisonously.  Blood 
containing  carbon  monoxide  is  characterized  by  very  distinct  spectrum 
reactions. 


In  1890  L.  Mond,  in  conjunction  with  other  chemists,  discovered  an 
exceedingly  remarkable  property  of  carbon  monoxide,  viz.,  that  it  com- 
bined with  very  finely  divided  nickel  at  25-30°  to  form  a liquid,  which 
volatilized  readily — nickel  carbonyl,  Ni(CO)^.  Since  then  it  has  been 
observed  that  carbon  monoxide  also  combines  with  other  metals  of  the 
iron  group,  forming  metal  carbonyls. 

Nickel  Carbonyl,  Ni(CO)4,  is  a colorless,  strongly  refracting  liquid, 
which  boils  at  43°  (751  mm.  pressure)  and  becomes  crystalline  at  — 25°. 

20 


234 


INORGANIC  CHEMISTRY. 


Its  vaj)C)rs  decom))ose  witli  ex])losi()n  at  6o°  ; in  the  air  they  burn  with  a 
very  smoky  flame  [see  Z.  f.  ])hys.  (ill.  8 (1891),  150]. 


Carbon  monoxide,  being  an  unsaturated  compound,  combines  like 
ethylene  (]).  154)  with  two  atoms  of  chlorine  to  Carbonyl  Chloride  or 
Phosgene  Gas,  COClj,  the  chloride  corresponding  to  carbonic  acid  : 

CO  4-  Cb  = COCl,. 

I vol.  I vol.  I vol. 

It  is  obtained  by  bringing  together  equal  volumes  of  carbon  monoxide 
and  chlorine  in  direct  sunlight  (hence  the  name  from  light,  and 
yevvdio,  I produce) ; also  when  the  gases  are  conducted  over  ignited  plati- 
num sponge  or  animal  charcoal.  It  can  also  be  made  by  conducting 
monoxide  into  antimony  pentachloride,  SbCl^. 

This  compound  is  important  in  the  color  industry,  and  other  con- 
venient methods  for  its  jireiiaration  may  be  found  by  consulting  Erd- 
mann, Ber.  26  (1893),  II,  1990. 

It  is  a colorless,  suffocating  gas,  the  density  of  which  agrees  with  the 
molecular  formula  COCb-  Phosgene  can  easily  be  liquefied  by  chilling ; it 
boils  at  -|-8°  and  has  the  specific  gravity  1.43.  Water  decomposes  it  into 
hydrogen  chloride  and  carbon  dioxide  : 

COCI2  -f  HjO  = COj  + 2HCI. 


Amido-derivatives. — The  following  are  amido-derivatives  of  carbonic  acid  which  are 
fully  treated  in  Organic  Chemistry  and  will  receive  but  mere  mention  here  : Carbamic 
Nil  NH 

acid,  (^0<Cqj4^»  carbamide  or  urea,  They  bear  the  same  relation  to 

carbonic  acid  that  sulphamic  acid  and  sulphamide  sustain  to  sulphuric  acid : 


NfT 

CO<oH^ 


NH 
SO  <-^  2 


CO< 


NH, 

NH, 


SO 


3 

2 


and  can  be  obtained  in  a corresponding  manner  from  carbonyl  chloride  and  ammonia. 
Analogous  bodies  are  derived  from  hydrazine — the  hydrazides,  e.  g.,  carbohy  dr  azide. 

Mention  may  also  be  made  of  nitrogen  carbonyl,  prepared 

by  Curtius  and  Heidenreich.  It  is  the  counterpart  of  phosgene  but  is  obtained  from 
hydrazoic  acid.  It  is  a very  explosive,  exceedingly  volatile,  crystalline  compound.  In 
this  body  also  the  group  N3  deports  itself  like  the  halogens.  [See  p.  133  and  Ber.  27 
(1894),  2684;  Jr.  prakt.  Ch.  52  (1895),  454.] 


COMPOUNDS  OF  CARBON  WITH  SULPHUR. 

Carbon  Bisulphide,  CS.^,  is  formed,  like  the  dioxide,  by  the  direct 
union  of  carbon  and  sulphur;  if  vapors  of  the  latter  are  led  over  ignited 
carbon  the  escaping  bisulphide  vajiors  are  condensed  in  a cooled  receiver. 
Pure  carbon  bisulphide  is  a colorless,  mobile  liquid,  of  faintly  ethereal,  and 
in  no  sense  disagreeable,  odor;  it  solidifies  at  — 116°,  and  refracts  light 


CYANOGEN  COMPOUNDS. 


235 


strongly.  Its  specific  gravity  equals  1.29  at  0°.  It  is  very  volatile,  boils 
at  47°,  and  burns  with  a blue  flame  to  carbon  dioxide  and  sulphur  dioxide. 
When  a mixture  of  carbon  bisulphide  vapors  and  oxygen  is  ignited,  a 
violent  explosion  ensues : 

CS,  + 3O,  = CO,  + 2SO,. 

I vol.  3 V'ols.  I vol.  2 vols. 

In  nitric  oxide,  the  vapors  burn  with  an  intensely  brilliant  flame.  On 
directing  a strong  current  of  air  upon  carbon  bisulphide  in  a porcelain 
capsule  (which  conducts  heat  poorly),  so  much  heat  is  absorbed  by  the 
evaporation  that  the  residual  liquid  solidifies  to  a white,  snow-like  mass 
which  contains  water.  Carbon  bisulphide  is  insoluble  in  water ; but  mixes, 
in  every  proportion,  with  alcohol  and  ether.  It  dissolves  iodine  with  a 
violet-red  color,  and  serves  as  an  excellent  solvent  for  sulphur,  phosphorus, 
caoutchouc,  and  the  fatty  oils.  On  conducting  carbon  bisulphide  vapors 
over  heated  zinc-dust,  all  the  sul})hur  unites  with  the  zinc,  forming  zinc 
sulphide,  while  the  carbon  separates  as  soot: 

CS,  + 2Zn  = 2ZnS  -(-  C. 

Most  metals  react  in  a similar  manner. 

Carbon  bisulphide  may  be  viewed  as  the  anhydride  of  sulpho carbonic 
acid,  H2CS3.  The  salts  of  this  acid  are  obtained  by  the  solution  of  car- 
bon bisulphide  in  alkaline  sulphides  (see  Sulpho-salts,  p.  223) : 

CS,  T KjS  = K^CSj. 

On  adding  hydrochloric  acid  to  the  solutions  of  these  salts  the  sulpho- 
carbonic  acid  separates  as  a reddish-brown  oil.  This  decomposes 
readily. 

The  sulphur  compound  corresponding  to  carbon  monoxide  is  not  known  : there  exists, 
however,  one  containing  both  oxygen  and  sulphur — carbon  oxysulphide,  COS.  This  sub- 
stance is  produced  (in  small  quantity,  because  it  decomposes  at  about  the  same  tempera- 
ture) when  a mixture  of  sulphur  vapors  and  carbon  monoxide  gas  is  passed  through  red- 
hot  tubes. 

It  is  most  readily  obtained  from  potassium  sulphocyanide,  CN.  SK  (see  Organic  Chem- 
istry) by  the  action  of  dilute  sulphuric  acid.  Carbon  oxysulphide  is  a colorless  gas,  with  an 
odor  reminding  one  of  hydrogen  sulphide.  It  is  present  in  some  sulphur  springs.  The  gas  is 
very  readily  inflammable  and  burns  with  a blue  flame  to  carbon  dioxide  and  sulphurous  acid. 

It  is  soluble  in  an  equal  volume  of  water,  decomposing  gradually  (more  rapidly  in  the 
presence  of  alkalies)  into  the  dioxide  and  hydrogen  sulphide  : 

COS  -f  11,0  = CO,  + HS,. 


CYANOGEN  COMPOUNDS. 

Of  the  compounds  of  carbon  mention  will  also  be  made  of  those  of  cyanogen,  as  they 
are  of  importance  in  inorganic  chemistry. 

Nitrogenous  carbon  compounds  heated  with  potassium  hydroxide  yield  potassium  cya- 
nide, KCN,  which  with  ferrous  oxide  forms  the  so-called  yellow  prussiate  of  potassium, 
K^Fe(CN)p.  All  the  other  cyanogen  derivatives  may  be  prepared  from  these  two  com- 
pounds, They  all  contain  the  group  CN,  called  cyajtogcn.  It  is  similar  to  the  groups 
OH,  Nil,,  CH3,  and  is  a univalent  radical  (pp.  167,  171).  In  chemical  behavior  the 
cyanogen  group  is  very  similar  to  the  halogens  and  the  group  N3  of  hydrazoic  acid  ; with 


236 


INORGANIC  CHEMISTRY. 


tlie  metals  it  forms  metallic  cyanides  (KCN,  AgCN)  very  similar  to  the  haloid  salts. 
Hydrogen  cyanide  is  evolved  when  the  cyanides  are  heated  with  sulphuric  acid  : 

2KCN  1 HjSO^  = KjS(\  + 2IICN. 

Hydrogen  Cyanide,  IICN,  is  a colorless,  mobile  liejuid,  of  peculiar  odor,  and  boiling 
at  27°.  Like  the  halogen  hydrides,  it  is  an  acid,  forming  salts  with  metals  and  bases, 
and  is  known  as  hydrocyanic  or  prussic  acid.  Loth  it  and  its  salts  are  very  powerful 
poisons.  If  the  CN  group  is  separated  from  its  salts  it  doubles  itself,  yielding  dicy- 
anogen or  free  cyanogen,  CjNj  (N  C-C  N),  because,  like  the  other  univalent  groups 
(as  CII3,  see  p.  168),  it  cannot  exist  in  a free  condition. 


The  heats  formation  of  the  simplest  carbon  compounds  (from  the  diamond)  above  cited 
correspond  with  the  symbols  : 

(C,0)^-26.3  (C0,0)  = 68.3  (C,0,)  = 94.3  (CO^.Aq)  - 5-8. 

If  an  element  combine  with  another  according  to  multiple  proportions,  there  usually 
occurs,  in  the  union  of  the  first  atom,  a greater  disengagement  of  heat  than  with  the  fol- 
lowing atom.  The  numbers  above,  on  the  contrary,  show  that  the  union  of  the  second 
atom  of  oxygen  with  carbon  (CO,n)  sets  free  68.3  Cal.  ; that  of  the  first  atom  (C,0), 
however,  only  26.3  Cal.  This  can  only  be  explained  by  the  fact  that,  for  the  vapor- 
ization and  disaggregation  of  the  solid  carbon  molecules,  heat  is  necessary.  If  we 
assume  that  the  direct  union  of  the  first  atom  also  disengaged  68.3  Cal.,  it  would 
follow  from  this  that,  in  the  dissociation  of  12  parts  by  weight  of  carbon  into  gaseous  free 
atoms,  42.0  (—  68.3  — 26.3)  Cal.  were  absorbed. 


2.  OXYGEN  COMPOUNDS  OF  SILICON. 

Silicon  Dioxide,  Si02  {^Silicd),  the  only  well-known  oxide  of  silicon, 
is  widely  distributed  in  nature  as  rock-crystal,  quartz,  sand,  etc.  Quartz 
occurs  in  nature  crystallized  in  figures  of  the  hexagonal  system,  with  the 
specific  gravity,  2.6;  these  crystals  are  colorless,  or  colored  by  impurities. 

Quartz  is  a very  widely  distributed  mineral.  The  following  varieties,  distinguished  by 
color  and  general  appearance,  exist : Rock-crystal,  colorless  ; smoky  quartz,  brown  ; 
morion,  almost  black  ; citrine,  yellow  ; amethyst,  violet,  which  on  heating  becomes  yel- 
low. Common  quartz  is  cloudy,  slightly  transparent,  white  or  colored.  It  is  an  essen- 
tial constituent  of  granite,  gneiss,  quartz-porphyry,  sandstone  and  other  rocks.  Special 
varieties  of  it  are:  fibrous  quartz  (tiger’s-eye,  Africa),  the  cat’s-eye  (Ceylon),  etc.  The 
gray  hornstone,  the  green  chrysoprase  and  red  or  brown  jasper  consist  of  compact  quartz, 
etc. 

Tridyinite  (rptSv/wt,  triplets)  has  the  same  composition  as  quartz  and  crystallizes  in 
the  hexagonal  system.  It  is  found  in  volcanic  rocks.  Its  specific  gravity  is  2.3. 

Silica  may  be  obtained  in  an  amorphous  form  by  igniting  the  silicic 
acid  which  sei)arates  from  silicates,  but  in  fusions  it  has  been  obtained  in 
both  crystalline  forms.  The  natural  anhydride  always  contains  impuri- 
ties; the  purest  is  the  colorless  rock-crystal. 

Silicon  dioxide  is  insoluble  in  water  and  in  all  acids;  but  is  decom- 
l)osed  by  hydrolluoric  acid  with  the  formation  of  silicon  fluoride  (SiFlJ 
and  water  (j).  162).  Strong  ignition  with  sodium  or  ])otassium  reduces 
it  in  j)art  to  metallic  silicon.  The  reduction  is  readily  effected  by  the 


OXYGEN  COMPOUNDS  OF  SILICON. 


237 


use  of  metallic  magnesium  or  aluminium  [Gattermann,  Ber.  22  (1889), 
186;  Winkler,  ibid.  23  (1890),  2652;  Jahrb.  f.  Chem.  v (1895),  ^^]  • 

Si02  + 2Mg  = 2MgO  + Si. 

The  dioxide  prepared  artificially  dissolves  when  boiled  with  potassium  or 
sodium  hydroxide  ; the  natural  only  when  finely  divided  [Jahrb.  f.  Chem. 
VII  (1897),  85].  By  fusion  with  the  hydroxides  or  carbonates  of  the  alka- 
lies all  varieties  of  silica  yield  a glassy  mass  (water-glass)  soluble  in  water 
and  containing  silicates  (K^SiO^  or  KjSiOg).  Upon  the  addition  of 
hydrochloric  acid  to  the  aqueous  solution  of  the  potassium  or  sodium  salt, 
a very  voluminous,  gelatinous  mass  separates;  this  is  probably  norjnal 
or  orihosilicic  acid,  H^SiO^ : 

Na^SiO^  + 4HCI  = qNaCl  + H^SiO^. 

Orthosilicic  acid  is  also  formed  in  the  decomposition  of  silicon  fluoride 
by  water  (p.  163).  It  becomes  a white  amorphous  powder  having  the 
composition  H2Si03  or  3Si02.2H20  when  washed  with  water  and  dried 
in  the  air.  By  strong  ignition  water  is  gradually  expelled  until  finally 
the  anhydride  remains.  The  freshly  precipitated  acid  is  somewhat  solu- 
ble in  water,  more  readily  in  dilute  hydrochloric  acid  and  in  caustic  soda. 
On  adding  a solution  of  sodium  silicate  to  an  excess  of  dilute  hydro- 
chloric acid  the  silicic  acid  remains  dissolved.  From  the  hydro- 
chloric acid  and  sodium  chloride  solution  we  can  obtain  a perfectly  pure 
aqueous  solution  of  silicic  acid  by  dialysis  by  proceeding  in  the  fol- 
lowing manner : 

Pour  the  hydrochloric  acid  solution  into  a wide  cylindrical  vessel,  a.  a.  whose  lower  open- 
ing is  covered  with  animal  bladder  or  parchment  paper,  and  then  suspend  the  vessel  (dialy- 
ser)  in  another,  b,  containing  pure  water  (Fig.  67).  Osmosis  now  sets  in.  The  sodium 


Fig.  67. 


chloride  and  hydrochloric  acid  pass  through  the  parchment  paper  into  the  outer  water,  b, 
while  on  the  other  hand,  water  passes  from  the  outer  vessel  into  the  dialyser  ; the 
parchment  paper  is  not  permeable  to  silicic  acid.  This  alternate  diffusion  of  the  different 
particles  occurs  until  the  outer  and  inner  liquids  show  the  same  quantity  of  diffusible  sub- 
stances. Upon  introducing  the  dialyser  into  a fresh  portion  of  water,  the  dialysis  com- 
mences anew.  Finally,  after  repeated  renewal  of  the  external  water,  the  dialyser  will 
contain  a perfectly  pure  silicic  acid  solution,  free  from  sodium  chloride  and  hydrochloric 
acid.  The  solution  may  be  concentrated  by  evaporation  ; it  then  readily  passes  into  a 


238  INORGANIC  CHEMISTRY. 

gelatinous  mass.  1 lie  same  occurs  inslantancously  in  dilute  .solutions  if  a trace  of  sorlium 
carbonate  be  added  or  carbon  dioxide  be  led  into  it. 

Like  sodium  chloride,  all  cry.stallizable  soluble  substances  diffu.se  through  jiarchment. 
d'hese  are  known  as  crystalloids,  to  distinguish  them  from  the  non-dilfusible  colloids. 
To  the  latter  belong  gum,  gelatine,  albumin,  starch,  glue  (koaAo,  hence  the  name  col- 
loid), and  especially  most  of  the  sub.stances  which  occur  chiefly  in  vegetable  and  animal 
organisms.  Like  silicic  acid  these  colloids  exist  in  lifjuid,  soluble,  and  .solid  gelatinous 
conditions.  Many  other  substances  (like  ferric  and  aluminium  oxides)  which  ordinarily 
are  insoluble  in  water,  can  be  brought  into  acjueous  solution  by  dialysis. 

For  “osmotic  pressure”  and  the  determination  of  molecular  weights 
see  Solutions. 

We  have  already  seen  that  acids  like  sulphuric,  jthosphoric,  and  arsenic, 
are  capable  of  forming  anhydro-  or  poly-acids  by  the  union  of  several 
molecules  and  the  elimination  of  water  (j).  214).  Silicic  acid  is  particu- 
larly inclined  to  that  form  of  condensation.  It  forms  a large  number 
of  poly-silicic  acids,  Si./)3(OH)2,  Si30^(0H)^,  Si^CJ^OH),,  etc.,  derived 
from  the  normal  and  ordinary  acid,  according  to  the  common  formula: 

;;/Si(OII)^  — 

These  poly-acids  are  not  known  free  ; it  appears,  however,  that  many 
amorphous  forms  of  silica  occurring  in  nature,  as  agate,  chalcedony, 
opal,  which  lose  5-15  i)er  cent,  of  water  by  ignition,  represent  such  poly- 
acids. The  natural  silicates  are  the  salts  of  such  acids.  Agate  consists 
of  numerous  layers  of  variously  colored  chalcedony,  quartz  and 
amethyst.  The  majority  are  derived  from  the  acids  H2Si205,  H^SigOg, 
H2Si30^,  H^Si^Ojo,  and  others.  Only  a few  silicates  are  obtained  from  the 
normal  acid,  e.  g.,  chrysolite,  Mg2Si04. 

Corresponding  to  carbon  bisulphide,  CS2,  is 

Silicon  Disulphide,  SiS2,  which  may  be  made  by  heating  amorphous  silicon  with 
sulphur,  or  by  conducting  sulphur  vapors  over  an  ignited  mass  of  silica  and  carbon.  It  sub- 
limes in  shining,  silky  needles,  which  water  changes  to  silicic  acid  and  hydrogen  sulphide. 


Germanium,  Tin  and  Lead  belong  to  the  group  of  carbon  and  sili- 
con. They  yield  the  derivatives  MeOg  and  Me(OH)^  (pp.  226,  227),  but 
as  they  form  lower  oxides,  GeO  and  SnO,  which  manifest  a perfectly 
basic  nature  and  unite  with  acids  to  form  salts,  and  as  these  elements  are 
thereby  allied  to  the  metals,  they  will  be  discussed  under  that  division. 


TITANIUM.  ZIRCONIUM.  THORIUM. 

Ti  = 48.1.  Zr  = 90.6.  Th  = 232. 

The  same  relation  that  vanadium,  niobium,  and  tantalum  show  to  the  elements  of  the 
I)ho.sj)horus  group  is  manifested  by  the  three  elements  titanium,  zirconium  and  thorium 
for  the  silicon  group  (.see  Periodic  System,  p.  246). 

P = 31.0  V = 51.2  Si  = 28.4  Ti  = 48.1 

As  — 75  Nb  = 94  (le  =72  Zr  = 90.6 

Sb  120  Ta  = 183  Sn  = 118.5  Th  = 232. 


TITANIUM — ZIRCONIUM. 


239 


In  all  their  deportments  they  strongly  resemble  tin  ; they  possess,  however,  a more  metal- 
lic character  in  their  derivatives.  They  are  quadrivalent,  affording  compounds  of  the 
form  MeX^,  in  which  X represents  univalent  elements  and  groups  ; those  of  the  form 
MeX2,  corresponding  to  the  stannous  derivatives,  are  unknown.  The  hydroxides, 
Me(OH)^  and  MeO(OHj2,  have  a stronger  basic  nature  than  stannic  acid  and  form  stable 
salts  with  acids  ; the  basicity  increases  successively  with  the  atomic  weights.  Corre- 
sponding to  this,  the  acidity  of  the  hydrates,  i.  e.,  their  capability  of  exchanging  hydrogen 
for  metals,  gradually  diminishes.  Thorium  hydroxide,  Th(OH)^,  is  not  able  to  form 
metallic  salts. 


TITANIUM. 

Ti  = 48.1. 

This  metal  occurs  in  nature  as  titanium  dioxide  (rutile,  anatase,  brookite)  and  in  titan- 
ates  (perofskite,  CaTiOg,  menaccanite,  FeTiOg).  Free  titanium  is  a gray,  metallic 
powder,  obtained  by  Berzelius  on  heating  potassium  titanium  fluoride  (K2TiFlg)  with 
potassium.  Wohler,  however,  was  the  first  to  recognize  it  as  metal.  It  burns  when 
heated  in  the  air,  and  decomposes  water  on  boiling.  It  dissolves  in  dilute  hydrochloric 
and  sulphuric  acids,  with  evolution  of  hydrogen. 

Titanium  Chloride,  TiCl^,  is  formed,  like  silicon  chloride,  by  conducting  chlorine 
over  an  ignited  mixture  of  the  dioxide  and  carbon.  A colorless  liquid,  of  specific  gravity 
1.76,  fuming  strongly  in  the  air  (with  decomposition  into  hydrochloric  and  titanic  acids), 
and  boiling  at  136°.  It  solidifies  at  — 25°.  The  vapor  density  corresponds  to  the  molec- 
ular formula  TiCl^.  It  behaves  like  tin  tetrachloride  with  water.  A compound — hydro- 
fliiotitanic  acid,  HgTiP’g — and  its  salts  are  known.  They  correspond  to  hydrofluosilicic 
acid. 

Titanic  Acid,  H^TiO^  (first  obtained  pure  by  H.  Rose),  separates  as  a white,  amor- 
phous powder,  on  adding  ammonium  hydroxide  to  the  hydrochloric  acid  solution  of  the 
titanates.  When  dried  over  sulphuric  acid  it  loses  i molecule  of  water  and  becomes 
TiO(OH)2.  Titanic  acid,  like  silicic  and  stannic  acids,  forms  poly-acids.  The  hydrates 
dissolve  in  alkalies  and  strong  acids,  to  form  salts.  On  igniting  the  hydroxides  we  get 

Titanium  Dioxide,  Ti02,  which  may  be  procured  crystallized  as  rutile,  brookite,  and 
anatase.  When  ignited  in  a stream  of  hydrogen  it  changes  to  the  oxide  '1^20,.  Titanium 
dioxide  is  almost  insoluble  in  acids  ; it  is  only  dissolved  by  hydrofluoric  acid.  It  forms 
titanates  upon  fusion  with  the  alkalies. 

The  hydroxides,  TiO^H^,  Ti03H2,  etc.,  conduct  themselves  as  feeble  bases  with  strong 
acids  and  afford  salts  with  them  (<?.  g.,  TiO.  SO4),  which  are  decomposed  by  water.  The 
alkaline  titanates  (K2TiOg)  are  very  unstable.  Other  titanates  occur  in  nature,  e.  g., 
CaTiOg,  MgTiOg,  and  the  so-called  titanic  iron,  FeTiOg. 

Titanium  also  forms  derivatives  after  the  types,  Ti.^Og  and  TiO,  c.  g.,  TigClg  and  Ti2Cl4. 
The  sesquioxide  compounds  are  usually  green  or  violet  in  color,  while  those  of  the  mon- 
oxide form  are  black  or  brown. 

Titanium  yields  various  compounds  with  nitrogen.  When  the  dioxide  is  heated  in 
ammonia  gas,  a dark-violet  powder  of  the  composition  TiN2  results.  The  compound 
TigCN^ — the  so-called  cyan-titanium  nitride — is  sometimes  found  in  copper-red,  metallic 
cubes,  in  blast-furnace  slag,  when  iron  ores,  containing  titanium,  have  been  fused. 


ZIRCONIUM. 

Zr  = 90.6. 

Zirconium  is  very  rare  in  nature,  but  is  generally  found  in  silicates,  and  especially  as 
zircon,  ZrSiO^.  Free  zirconium,  like  carbon,  is  known  in  three  allotropic  forms.  On 
fusing  potassium  zirconium  fluoride  with  aluminium  L.  Troost  obtained  it  in  brilliant, 
hard,  steel-gray,  fiat  leaflets,  of  specific  gravity  4.15,  which  burned  in  the  oxyhydrogen 
flame  to  the  dioxide.  L.  Troost  also  i.solated  the  graphitic  form.  The  amorphous  variety 
was  observed  by  Berzelius,  who  was  the  first  chemist  after  Klaproth,  its  discoverer,  to 
engage  in  a study  of  zirconium.  lie  obtained  it  by  the  method  used  for  titanium  and 
got  it  as  a black  powder  which  burnt  readily.  Zirconium  tetrachloride,  ZrCl^,  and  fluoride, 
ZrFl^,  are  very  similar  to  the  corre.sponding  titanium  compounds. 

Zirconic  acid,  Zr(OH)^,  is  not  definitely  known.  Metazirconic  acid,  ZrO(OH)2,  is 


240 


INORGANIC  CHEMISTRY. 


[)recipitatecl  on  adding  ammonia  water  to  acid  solutions,  in  the  form  of  a white,  volumi- 
nous mass,  which  yields  zirconium  dioxide,  Zr(  ).^,  u|)on  ignition.  A colorless,  transparent 
crystalline  modification  is  known — isomorphous  with  rutile  and  cassiterite.  Metazircotiic 
acid  is  insoluble  in  caustic  potash  and  soda  ; it  only  forms  zirconates,  Na^Zrt ).,  and  Na^Zrt 
when  fused  with  alkalies  and  alkaline  carbonates.  Water  decomjjoses  tiiese  salts  witii 
the  liberation  of  alkali,  h'reshly  precipitated  metazirconic  acid,  ZrO(On)2,  dissolves 
readily  in  strong  acids.  Zirconium  dioxide  is  soluble  in  hot,  concentrated  sulphuric  acid, 
forming  Zr(S()^)2,  which  can  also  be  obtained  from  a(jueous  solution  in  crystals, 

Zr(SOj2.4lI.p.  . 

In  late  years  zirconium  dioxide  has  become  of  great  imi)ortance  for  illuminating  pur- 
poses. It  is,  therefore,  i)rei)ared  in  large  amounts  from  zircon.  'I'he  finely  pulverized 
mineral  is  mixed  with  carbon  and  while  being  gently  heated  chlorine  is  conducted  over 
the  mixture  when  silicon  chloride,  SiCl^,  and  zirconium  chloride,  ZrCl^,  are  produced  ; or 
the  mineral  is  decomposed  with  hydrofluoric  acid.  IJerzelius  noticed  the  light-emission 
power  of  the  oxide  when  it  was  heated.  Caron  (1868)  determined  its  value  for  the  tem- 
perature of  the  oxyhydrogen  flame  and  d'essie  du  Mothay  endeavored  to  utilize  it  practi- 
cally. Linnemann  (1885)  made  the  oxide  into  thin  plates  and  soon  thereafter  W.  Kochs 
succeeded  by  a peculiar  process  in  shaping  it  into  any  desired  form,  d'he  durability  of 
the  zirconia  light-giving  bodies  in  the  oxyhydrogen  flame,  the  fact  that  they  do  not  decom- 
pose in  the  air,  that  the  light  which  they  diffuse  is  intensely  white  (its  continuous  spec- 
trum includes  the  P'raunhofer  lines  from  A to  H)  render  it  superior  to  the  Drummond 
lime  light.  It  is  extensively  used  for  projection  lamps,  for  photography  and  in  optical 
researches. 

THORIUM. 

Th  = 232. 

Until  recently  thorium  was  only  known  as  a constituent  of  certain  rare  northern  min- 
erals. Berzelius  first  discovered  it  in  the  thorite  from  Arendal,  later  Wohler  found  it  in 
pyrochlore,  and  Kersten  in  monazite.  Small  amounts  of  it  are  present  in  many  orthites. 
Its  application  to  illuminating  purposes  has  created  a demand  for  large  quantities  of  it 
and  at  present  it  is  isolated  from  a rich  monazite  deposit  in  McDowell  County,  N.  C. 
That  mineral  contains  5-6  per  cent,  of  thorium  oxide. 

Free  thorium  has  been  obtained  as  a gray,  crystalline  powder  of  specific  gravity  i i.o  by 
reduciirg  potassium  thorium  fluoride  or  potassium  thorium  chloride  with  metallic  sodium 
or  potassium.  Its  specific  heat  is  0.0276.  It  burns,  when  heated  in  the  air,  to  dioxide, 
thoria,  Th02.  The  metal  is  easily  soluble  in  concentrated  sulphuric  acid,  nitric  acid,  and 
hydrochloric  acid.  It  is  not  soluble  in  the  alkalies.  The  tetrachloride,  ThCl^,  results 
from  the  action  of  hydrogen  chloride  upon  metallic  thorium  ; it  melts  at  a white  heat  and 
sublimes  in  white  needles.  Its  vapor  density  corresponds  to  the  formula  ThCl^.  It  is 
soluble  in  water  and  separates  in  crystals  of  the  formula  ThCl^  . 8H2O.  It  forms  crystal- 
line double  salts  with  the  alkaline  chlorides.  The  fluoride,  ThFl^,  is  a white  powder, 
insoluble  in  water.  The  hydroxide,  Th(0H)4,  is  precipitated  as  a white  jelly  on  adding 
ammonia  water  to  thorium  salts.  The  hydroxide,  ThO(OH)2,  is  not  known.  On  heat- 
ing thorium  hydroxide,  Th (011)4,  the  dioxide,  ThOg,  is  obtained.  The  latter  is  not 
soluble  in  dilute  acids  ; it  is,  however,  converted  into  thorium  sulphate,  Th(S04)2,  when 
acted  upon  with  concentrated  sulphuric  acid.  The  sulphate  dissolves  readily  in  cold 
water  and  crystallizes  from  the  same  in  colorless,  brilliant  needles  of  the  formula 
'141(804)2. When  its  aqueous  solution  is  heated  almost  all  of  the  sulphate  sepa- 
rates as  a woolly  mass,  which  on  cooling,  slowly  dissolves.  The  hydrate  and  dioxide  do 
not  form  .salts  with  the  alkalies.  The  i.somorphism  of  thorium  dioxide  with  uranium 
dioxide,  and  thorium  .sulphate  with  uranium  sulphate  is  very  interesting. 

Thorium  dioxide  is  even  more  valuable  for  lighting  purposes  than  zirconium  dioxide 
because  it  gives  forth  a blinding  light  at  the  temperature  of  the  Bun.sen  flame — light  rays 
of  every  refrangibility  are  emitted.  Auer  von  Websbach  has  utilized  this  property  in  a 
gas-light.  Fine-meshed  material  of  cotton  or  linen  .saturated  with  a .solution  of  thorium 
nitrate  and  a little  cerium  nitrate,  surrounds  the  colorless  flame  of  a Bunsen  burner. 
When  the  vegetable  material  has  burned  away  a mantel  of  thorium  dioxide  and  cerium 
oxide  remains,  which  has  the  exact  form  of  the  original  substance,  and  emits  a bright, 
clear,  soft  light.  'I'he  ordinary  “ gh)w-mantle.s  ” contain  about  98-99  percent  of  thoria 


BORON. 


241 


and  2-1  per  cent,  cerium  oxide  ; it  is  exactly  this  mixture  which  affords  the  desired  white 
light.  The  intense  emission  power  Bunte  believes  is  due  to  the  catalytic  action  of  finely 
divided  cerium  oxide,  which  causes  a rapid  combustion  and  consequently  a high  tem- 
perature. Hence  in  the  Welsbach  mantle  the  thorium  oxide  simply  acts  as  a stable, 
porous  carrier  for  the  cerium  oxide. 


BORON. 

B = II. 

This  element  is  generally  classed  with  the  metalloids,  and  stands 
isolated  among  them ; it  forms  the  transition  from  these  to  the  metals, 
which  is  manifest  from  its  position  in  the  periodic  system.  On  the  one 
side,  especially  when  free,  it  resembles  carbon  and  silicon  ; on  the  other, 
it  approaches  the  metals,  beryllium,  aluminium,  and  scandium  (see  the 
Periodic  System  of  the  Elements).  As  recently  observed,  it  affords  a 
gaseous  hydride,  but  it  is  not  very  stable,  and,  like  stibine,  may  be  easily 
decomposed  into  its  constituents.  Its  oxide,  B.^03,  although  really  of  an 
acid  nature,  approaches  such  metallic  oxides  as  aluminium  oxide,  AI2O3, 
which  functionates  both  as  base  and  acid.  Boron  is  trivalent,  and  yields 
only  compounds  of  the  form  BX3. 

It  is  found  in  nature  as  boracic  acid  and  in  the  form  of  borates,  like 
borax  (sodium  salt),  boracite  (2Mg3BgOj5  -)-  MgCl2 ; Stassfurt).  Boron 
exists  in  two  allotropic  modifications:  amorphous  and  crystalline. 

Mois.san  made  aynorphotts  boron  by  igniting  a mixture  of  freshly  fused  boric  acid  with 
magnesium  powder,  free  from  iron.  By  a rather  circuitous  treatment  of  the  fusion  the 
boron  (with  2 per  cent,  impurities)  is  obtained  as  a black  powder  of  specific  gravity  2.45. 
The  product  obtained  heretofore  by  other  methods,  and  believed  to  be  amorphous  boron, 
was  a mixture  of  boron,  iron  boride,  boron  nitride,  etc.  The  boron  of  Moissan  ignites 
at  700°  in  the  air  ; at  higher  temperatures  sulphur,  chlorine,  bromine,  nitrogen,  silver 
and  platinum  combine  readily  with  it.  It  has  a great  affinity  for  oxygen,  hence  acts 
as  a powerful  reducing  agent.  When  rubbed  together  with  lead  dioxide  explosion 
occurs.  Oxygen  acids  are  comparatively  easily  reduced  by  boron  ; potassium  perman- 
ganate, silver  nitrate  and  ferric  chloride  suffer  this  change  in  the  cold. 

The  crystalline  variety  may  be  obtained  by  igniting  boron  trioxide  with  aluminium. 
The  boron,  separated  by  the  aluminium,  dissolves  in  the  excess  of  the  latter,  and  crystallizes 
from  it  on  cooling  ; upon  dissolving  the  aluminium  in  hydrochloric  acid  the  boron  remains 
in  shining,  transparent,  quadratic  crystals,  which  are  more  or  less  colored,  and  have  a 
specific  gravity  of  2.63.  The  crystals  are  not  pure  boron,  but  contain  aluminium  and  car- 
bon. In  their  luster,  refraction  of  light,  and  hardness,  they  resemble  the  diamond. 
Cry.stalline  boron  is  more  stable  than  the  amorphous  ; it  does  not  oxidize  upon  ignition, 
and  is  only  slightly  attacked  by  acids.  Fused  with  potassium  and  sodium  hydroxide  both 
modifications  yield  borates. 


Boron  Hydride,  probably  BH3,  is  only  known  mixed  with  hydrogen.  It  results 
when  hydrochloric  acid  acts  upon  magne.sium  boride.  A solid  boron  hydride  is  produced 
on  heating  boron  trioxide  with  sodium,  or  borax  with  magnesium  pow^der.  It  has  not 
been  obtained  pure. 

Boron  Trichloride,  BCI3,  may  be  prepared  by  heating  boron  in 
chlorine,  or  conducting  a stream  of  the  latter  over  an  ignited  mixture  of 
the  trioxide  and  carbon  (see  Silicon  and  Aluminium  Chlorides)  : 

+ 3C  + — -BCL  -P  3CO. 


21 


242 


inor(;anic  chemistry. 


It  is  a colorless  liquid,  of  specific  gravity  1.35,  and  boiling  at  17°.  Its 
vapor  density  corresponds  to  the  molecular  formula  BCl,.  The  liquid 
fumes  strongly  in  the  air  and  is  decomposed  by  water  into  boric  and 
hydrochloric  acids : 

I3CI3  + 311,0  = B(OII)3-f3lICl. 

The  trichloride  also  results  from  the  action  of  the  pentachloride  of 
phosphorus  upon  boron  trioxide  : 

b-Ps  + 3bCb  = 2BCI3  + 3POCI3. 

Boron  Fluoride,  BFI3,  is  similar  to  silicon  fluoride,  and  is  produced 
according  to  the  same  methods,  by  the  action  of  hydrofluoric  acid  upon 
the  trioxide,  or  by  warming  a mixture  of  the  trioxide  and  calcium  fluo- 
ride with  sulphuric  acid  : 

B,03  + 3CaFl,  + 3lI,SO,  = 3CaSO,  + 311,0  + 2BFI3. 

It  is  a colorless  gas,  fuming  strongly  in  the  air,  of  specific  gravity  66 
(O,  = 32),  and  may  be  condensed  to  a liquid  under  strong  pressure.  It 
dissolves  very  readily  in  water  (700  volumes  in  i volume),  producing 
Hydrogen  Borofluoride,  BFl^H  (::r=  BFI3.  HFl),  which  remains  in 
solution  : 

4BFI3  -f  311,0  = 3HBFh  + H3BO3. 

The  reaction  is  analogous  to  the  formation  of  hydrofluosilicic  acid 
from  silicon  fluoride  (see  p.  163).  Hydrogen  borofluoride  is  a monobasic 
acid,  only  known  in  its  solution  and  in  its  salts. 

Boric  Acid,  H3BO3  = B(OH)3,  occurs  free  in  nature  and  in 
salts.  In  some  volcanic  districts,  especially  in  Tuscany,  the  steam 
escaping  from  the  earth  (fumeroles,  etc.)  contains  small  quantities 
of  it.  These  vapors  condense  in  small  natural  pools,  or  are  conducted 
into  walled  basins.  By  concentration  of  the  aqueous  solution,  boric 
acid  separates.  To  prepare  pure  boric  acid,  precipitate  a hot  solu- 
tion of  borax  with  nitric  acid.  The  acid  seiiarates  in  colorless,  shining 
scales  ; it  dissolves  in  25  parts  of  water  at  14°,  or  in  3 parts  at  100°.  The 
solution  shows  a feeble  acid  reaction  with  litmus;  turmeric  paper, 
moistened  with  it,  is  colored  red-brown  after  drying.  On  boiling  the 
solution,  boric  acid  escapes  with  the  steam.  An  alcoholic  solution 
of  the  acid  burns  with  a green  flame.  These  reactions  afford  a ready 
means  for  its  detection. 

When  heated  to  100°,  the  acid  loses  one  molecule  of  water,  and  passes 
into  the  anhydro-  or  meta-acid,  HBO,,  which  at  140°  is  converted  into 
tetraboric  acid,  H,B^O^.  When  ignited,  boric  anhydride  or  Boron  tri- 
oxide, B.,03,  is  produced.  This  is  a fusible,  glassy  mass,  of  specific 
gravity  1.8,  and  is  slightly  volatile  at  a very  high  heat.  Water  dissolves 
the  anhydride  to  boric  acid. 

It  is  a very  weak  acid  ; and  can  be  ex])clled  from  its  salts  by  most 
other  acids.  By  fusion  it  removes  most  acids  from  their  salts,  in  con- 
sc(|uence  of  the  difficult  volatility  of  its  anhydride. 

Salts  of  normal  boric  acid,  B(OII)3,  are  not  known,  while  the  ethers, 


PERIODIC  SYSTEM  OF  THE  ELEMENTS. 


243 


B(0.  CH3)3,  are.  The  salts  of  metaboric  acid,  BO . OK,  can  be  obtained 
crystallized,  but  they  are  very  unstable.  They  are  decomposed  by  car- 
bon dioxide  with  production  of  salts  of  tetraboric  acid  : 

4NaB02  -b  CO2  = Na^B.O^  + Na2C03. 

The  latter,  from  which  the  ordinary  borates  are  derived  (see  Borax), 
may  be  viewed  as  an  anhydro-acid,  produced  by  the  union  of  four  mole- 
cules of  trihydric  boric  acid  (compare  p.  238): 

4B(0H)3-5H20  = H2B,0,. 

On  heating  amorphous  boron  in  a stream  of  nitrogen  or  ammonia,  or  by  igniting  a 
mixture  of  the  trioxide  and  carbon  in  nitrogen  gas,  there  is  formed  Boron  nitride^  BN. 
This  is  a white  amorphous  powder,  which  gives  forth  an  extremely  intense  greenish-white 
light  when  heated  in  a gas  flame.  Boric  acid  and  ammonia  result  when  steam  at  200° 
is  conducted  over  the  nitride  : 

BN  + 3H2O  = B(0H)3  4-  NH3. 

In  the  heat  of  the  electric  furnace  carbon  unites  with  boron  to  Boron  carbide,  BgC, 
which  is  very  similar  to  carborundum,  but  is  harder  than  the  latter.  It  consists  of  black, 
brilliant  crystals,  of  specific  gravity  2.5.  It  is  exceedingly  stable  (Moissan). 


PERIODIC  SYSTEM  OF  THE  ELEMENTS. 

In  the  preceding  pages  we  have  studied  the  non  metals  and  their  com- 
pounds with  hydrogen,  the  halogens  and  oxygen.  In  their  entire 
behavior  they  arrange  themselves  in  four  natural  groups,  the  members  of 
which  show  unmistakable  family  similarities.  It  is  only  in  the  case  of 
boron  and  of  hydrogen  and  the  new  constituents  of  air  that  allied  mem- 
bers have  not  been  found.  Attention  has  been  repeatedly  called  to  the 
gradation  in  the  similarity  of  the  group  members,  which  seems  to  increase 
with  the  rise  of  the  atomic  weight,  and  also  to  the  relations  of  the  various 
groups  to  one  another — relations  observed  among  the  metals.  All  these 
points  appear  more  striking  and  more  regularly,  if  the  elements  be  con- 
sidered in  the  arrangement  which  the  periodic  system  assigns  them. 

From  the  time  the  first  atomic  weight  determinations  were  made,  regu- 
larities between  the  atomic  weights  and  the  properties  of  the  elements 
were  thought  to  exist.  One  of  the  first  chemists  to  surmise  this  was  J. 
W.  Ddbereiner,  whose  views  were  clear  and  definite  in  contrast  to  the 
fantastic,  mathematical  sophistries  of  many  of  his  successors,  especially 
those  of  the  Frenchman  Chancourtois,  who  recently  has  been  unjustly 
heralded  as  the  forerunner  of  D.  Mendelejeff  and  Lothar  Meyer,  the 
founders  of  the  periodic  system.* 


*See  Dobereiner,  Pogg.  Ann.  15  (1829),  301  ; Lothar  Meyer,  Ann.  Chem.  Suppl. 
7 (1871),  354;  Mendelejeff,  ibid.,  Supjjl.  8 (1872),  133;  re])rinted  in  No.s.  66  and 
68  of  Ostwald’s  Klas.siker  der  exakten  Wissen.schaften.  For  tlie  history  of  tlie  sys- 
tem see  Mendelejeff,  Princiifles  of  Chemistry  (1891),  683,  692;  L.  Meyer,  Modern 
Theories  of  Chemistry  (1883) ; also  K.  Seubert,  Z.  f.  anorg.  Chem.  9 (1895),  334. 


244 


INORGANIC  CHEMISTRY. 


Arranging  the  elements  according  to  increasing  atomic  weight  we 
observe  that  similar  elements  return  after  definite  intervals,  'rims  they 
arrange  themselves  in  several  })eriods: 


I. 

I.i 

P)e 

P C 

N 

0 

LI 

2. 

Na 

Mg 

A1  Si 

P 

S 

Cl 

3- 

K 

Ca 

Sc 

Ti  V 

Cr 

Mm 

be 

Ni  Co 

Cu 

Zn 

Ca  (ie 

As 

Se 

4- 

Rb 

Sr 

Y 

Zr  Nb 

Mo 

— 

Ru 

Rh  Pd 

Ag 

Cd 

1 11  Sn 

Sb 

d'e 

5- 

Cs 

Ba 

La 

(Ce  Nd 

Pr) 

— 

— 

6. 

— 

— 

Yb 

— Ta 

W 

— 

Os 

Ir  Pt 

All 

Hg 

Tl  Pb 

Pi 



7- 

Hydrogen  has  no  family;  it  alone  would  form  the  first  period,  and  is 
therefore  omitted,  as  are  also  helium,  argon,  and  the  other  recently  dis- 
covered constituents  of  the  air,  because  their  chemical  nature  is  not 
known;  the  positions  of  tellurium  and  iodine  have  also  been  exchanged, 
on  the  presumption  that  the  atomic  weight  assigned  tellurium  at  present 
is  too  high. 

The  period  syste?n,  or  the  law  of  periodicity,  which  is  indicated  by  this 
arrangement,  declares  that  the  properties  of  the  ele^nents  are  periodic  func- 
tions of  their  atomic  weights. 

The  first  two  series,  lithium  (Li)  to  fluorine  (FI),  and  sodium  (Na)  to 
chlorine  (Cl),  present  two  periods  of  seven  members  each,  in  which  the 
corresponding  (above  and  below)  members  exhibit  a great  but  not  com- 
plete analogy.  Sodium  resembles  lithium ; magnesium,  beryllium;  chlo- 
rine, fluorine,  etc.  Then  follow  two  periods,  consisting  of  seventeen  ele- 
ments each  : potassium  (K)  to  bromine  (Br),  and  rubidium  (Rb)  to  iodine 
(I).  The  series  5 and  bare  incomplete,  and  together  probably  constitute  a 
period.  In  the  seventh  series  there  are  as  yet  but  two  elements  : thorium 
= 232  and  uranium  ==  239.5.  Thus  result  three  great  periods,  whose  cor- 
responding members  exhibit  an  almost  complete  analogy:  the  elements 
K,  Rb,  Cs;  Ca,  Sr,  Ba;  Ga,  In,  T1 ; As,  Sb,  Bi,  etc.,  are  so  similar  that 
they  remind  us  of  the  homologous  series  of  the  carbon  compounds  (com- 
pare p.  153),  and,  therefore,  can  be  designated  as  homologous  elements. 
It  is  only  in  the  third  great  period  (series  5 and  6)  that  the  middle 
members  exhibit  any  variations. 

Now  on  comparing  the  three  great  periods  with  the  two  small  ones,  we 
discover  that  the  first  members  are  analogous  to  each  other;  K,  Rb,  Cs, 
resemble  Na  and  la  ; Ca,  Sr,  Ba,  resemble  Mg  and  Be.  Then  the  simi- 
larity gradually  lessens,  disappears  apparently  in  the  middle  members, 
and  only  aj)j)ears  again  toward  the  end  of  the  periods:  I and  Br  resemble 
chlorine  and  fluorine;  Teand  Se,  sul|)hur  and  oxygen  ; Bi,  Sb,  As,  phos- 
phorus  and  nitrogen,  etc.  The  character  or  the  function  of  the  three 
great  periods  is  therefore  other  than  that  of  the  two  small  periods.  But 
in  all  five  periods  we  can  detect  a gradual,  regular  alteration  in  the  prop- 
erties of  the  adjoining  heterologous  This  is  particularly  mani- 

fest in  the  measurable  physical  jiroperties,  all  of  which  show  a maximum 
or  minimum  in  the  middle  of  the  periods  (of  both  the  great  and  the 
small ). 

'I'he  ‘^ame  regularity  exhibits  itself  even  in  chemical  properties,  in  the 


PERIODIC  SYSTEM  OF  THE  ELEMENTS. 


245 


two  small  ])eriods,  esj)ecially  in  the  valence  of  the  elements  in  their  com- 
pounds with  hydrogen  or  the  hydrocarbon  groups  CH.^,  etc.  (desig- 

nated by  R;  compare  ]).  247).  The  hydrogen  valence  rises  and  falls 
periodically  with  the  specific  gravity  : 

I II  III  IV  III  II  I 

NaR  MgR2  AIR3  Silb  bHg  SII^  CIH. 

On  the  other  hand,  the  maximum  valence  of  the  elements  with  refer- 
ence to  oxygen  increases  gradually  if  chlorine  be  regarded  as  septivalent 
in  perchloric  acid  (pp.  171,  173)  : 

I IT  III  IV  V VI  VII 

Na20  MgO  AI2O3  Si02  SO3 

The  chemical  valence  expresses  itself  somewhat  differently  in  the  three 
great  periods.  In  them  we  have  a double  periodicity  ; thus,  e.  g.,  with  the 
salt-forming  oxides:  * 

I II  III  IV  V VI  VII 

K2O  CaO  SC2O3  Ti02  V2O5  Cr03  Mn207 

II  II  III  IV  V VI  VII 

CuO  ZnO  Ga203  Ge02  AS2O5  (8003)  (Br20d. 

In  consequence  of  this  double  periodicity,  the  first  seven  and  the  last 
seven  members  of  the  two  great  periods,  with  respect  to  their  valence 
(and  consequently  also  their  compounds),  resemble  the  seven  members 
of  the  two  small  periods.  To  bring  out  this  double  periodicity  and 
analogy,  the  first  seven  and  last  seven  members  of  the  great  periods  are 
divided  into  two  series,  and  arranged  under  the  corresponding  seven 
members  of  the  small  periods. 

In  this  way  the  three  middle  members  of  the  great  periods  (which  are 
found  between  the  dotted  lines  of  the  table,  p.  244)  come  to  stand  apart, 
as  they  have  no  analogues.  In  this  manner  arises  the  following  table,  in 
which  the  seven  (or  ten)  vertical  columns  include  homologous  elements : 


Li 

Be 

B 

C 

N 

0 

FI 

Na 

Mg 

A1 

Si 

P 

S 

Cl 

K 

Ca 

Sc 

Ti 

V 

Cr 

Mn 

Fe  Ni  CO 

Cu 

Zn 

Ga 

1 Ge 

As 

Se 

Br 

Rb 

Sr 

Y 

Zr 

Nb 

Mo 

— 

Ru  Rh  Pd 

Cd 

In 

Sn 

Sb 

Te 

I 

Cs 

Ba 

La 

Ce 

(Nd,Pr) 

— 

— 

— 

— 

Yb 

— 

Ta 

W 

— 

Os  Ir  Pt 

Au 

Hg 

T1 

Bi 

Pb 

— 

— 

In  the  table  (on  p.  246)  we  have  presented  the  same  grouping  of  the 
elements,  together  with  their  atomic  weights,  given  in  round  numbers. 

When  the  periodic  grouping  of  the  elements  was  first  presented,  the  atomic  weight  of 
indium  and  that  of  uranium  had  to  be  doubled,  if  they  were  to  occupy  the  places  indi- 


*The  elements  are  given  here  with  their  highest  oxygen  valence.  Perbromic  anhy- 
dride, Pr20-,  and  selenic  anhydride,  SeO.,,  have  not  been  prepared,  although  some  of 
their  derivatives  are  known.  Iron,  cobalt  and  nickel  have  been  omitted  because  their 
highest  oxides  are  not  well  defined. 


The  Periodic  Systef?i  of  the  Elefne^its. 


246 


INORGANIC  CHEMISTRY 


VIII 

Group. 

ISIO 

On 

\r^ 

0 

u 

0 

tr> 

1 : 
p^ 

0" 

'Pa 

Ov 

VT) 

Ya 

Rh  103 

1 - 
}m 

0 

Pa 

MD 

vr> 

0 

Ph 

Ru  102 

Os  191 

\TI 

Group. 

On 

E 

Cl  35-4 

0 

CO 
vri  ^ 

0^ 

r 

1 1 

0 

\r% 

C^  1 

1 

\T 

Group. 

e’o” 

IS 

VO 

0 

N 

ro 

m 

vr> 

Mo  96 
(?)Te  127 

1 

1 T 

CO 

1 

Ov 

rn 

cs 

Li 

V ’ 

Group. 

ro 

pH 

t/) 

p 

0 

(S 

'i- 

CV^Q 
^ C/2 

Y 

00  _ 

r E w 

1 

IV 

Group. 

^6 

PP 

u 

00 

M 

’(h 

r) 

00 

H 

Zr  90 

Sn  1 18 

~"n  1 0 

0 

E ^ 

OJ  1 

U 1 

cs 

rn 

cs 

H 

1 

II  III 

Group.  Group. 

CO 

pq 

< 

0 

-tO 

0 

m 

Y 89 

In  1 14 

1 La  138 

Yb  173 

T1  204 

- '-1 

li 

Ov 

o 

pq 

'+ 

rt 

b/) 

S 

i 

Ca  40 

1 Zn  65 

Sr  87 

Cd  112 

Ba  137 

Hg  200 

- '-i 

I 

Group. 

iS* 

E 

ro 

N 

vO 

OVr^ 

ro^ 

E 

Rb  85 

Ag  108 

Cs  133 

Au  197 

1 

1 

H-Compounds. 
Highest  salt-forming 
oxides. 

t/i  ^ "5 -5  'B  'B'B'B 

.U  „ N tovO  00  0 

<v 
c/) 

t/5 

^ ^ ^ j>  P 

<u 

Oh 

1 

PERIODIC  SYSTEM  OF  THE  ELEMENTS. 


247 


cated  for  them  by  the  system.  All  such  alterations  have  been  confirmed  by  recent  inves- 
tigations. The  series  of  gold,  iridium,  platinum  and  osmium  did  not  acsord  with  the 
system,  which  rather  demanded  the  following  : Osmium,  iridium,  platinum  and  gold. 
This  requirement  was  satisfied  by  a redetermination  of  the  atomic  weights  of  these  ele- 
ments by  Seubert  [Ann.  Chem.  (1891)  261,  272].  Hence,  the  periodic  system  offers  a 
control  for  the  numbers  of  the  atomic  weight,  while  formerly  they  appeared  to  be  irregu- 
lar, and,  at  the  same  time,  accidental.  We  are  consequently  justified,  until  we  have 
more  evidence  to  the  contrary,  in  assuming  that  the  determinations  of  the  atomic  weight  of 
tellurium  have  placed  that  value  too  high.  If  it  should  finally  be  proved  to  be  greater 
than  that  of  iodine,  then  the  periodic  system  would  be  seriously  affected  in  its  founda- 
tion ; it  would  then  lose  its  claim  to  being  a natural  law — for  this  would  not  tolerate  an 
exception. 

Mendelejeff,  on  the  basis  of  the  periodic  system,  predicted  in  a manner  similar  to 
that  shown  by  Leverrier  in  his  precalculation  of  neptune,  the  existence  of  new,  not  yet 
known,  elements  which  correspond  to  unoccupied,  free  places  or  gaps  in  the  table. 
In  fact,  three  such  gaps  have  been  filled  by  the  discovery  of  gallium,  scandium,  and 
germanium ; their  properties  have  shown  themselves  to  be  perfectly  accordant  with 
those  deduced  from  the  periodic  system.  [It  may  be  stated  here  that  Dobereiner,  on 

the  basis  of  the  law  of  Triads, — Cl,  Br,  I ; Li,  Na,  K ; Ca,  Sr,  Ba,  etc.:  = Br, 

etc., — predicted  two  elements,  which  should  form  a triad  with  fluorine,  and  one  a 
complete  series  with  phosphorus  and  arsenic,  but  the  prediction  has  not  been  realized.] 
At  present,  only  the  first  homologue  of  manganese  (with  atomic  weight  of  about  100)  is 
wanting.  In  the  very  incomplete  series  of  period  V,  it  is  probable  that  the  poorly  in- 
vestigated elements  mentioned  on  p.  27  will  find  place  (see  p.  244). 

The  entire  character  of  a given  element  is  determined  to  a very  high 
degree  by  the  law  of  periodicity;  hence,  all  physical  and  chemical 
properties  of  the  same  are  influenced  by  its  position  in  the  system. 
These  relations  we  will  examine  more  closely  in  the  individual  groups  of 
the  metals,  and  here  confine  ourselves  to  a notice  of  some  general  rela- 
tions, and  the  connection  of  atomic  weight  with  the  chemical  valence  of 
the  elements. 

The  relation  of  metalloids  to  metals  is  shown  with  great  clearness  in 
the  periodic  system.  The  first  members  of  all  periods  (on  the  left  side) 
consist  of  electro-positive  metals,  forming  the  strongest  bases,  the  alka- 
lies— Cs,  Rb,  K,  Na,  Li,  and  metals  of  the  alkaline  earths — Ba,  Sr,  Ca, 
Mg,  and  Be.  The  basic  character  diminishes  successively,  in  the  follow- 
ing heterologous  members,  and  gradually  passes  over  into  the  electro- 
negative, acid-forming  character  of  the  metalloids  FI,  Cl,  Br,  I.  Here 
is  observed  that,  in  the  periods  following  each  other,  with  higher  atomic 
weights,  the  basic  metallic  character  constantly  exceeds  the  metalloidal. 
The  first  period  comprises  five  metalloids  (B,  C,  N,  O,  FI),  the  second 
only  four  (Si,  P,  S,  Cl),  the  fourth  and  fifth  periods  each  only  three  (or 
two)  metalloids  (As,  Se,  Br,  and  Sb,  Te,  I),  which,  at  the  same  time, 
become  less  negative.  With  the  metalloidal  nature  is  combined  the 
power  of  forming  volatile  hydrogen  compounds.  Similar  volatile  deriva- 
tives are  also  afforded  by  the  metalloids  with  the  univalent  hydrocarbon 
groups  (as  CH3,  CjH^,  CgH^,  etc.),  which  resemble  hydrogen  in  many 
respects.  Such  metallo-organic  compounds,  in  which  the  elements  show 
the  same  valence  as  in  the  hydrogen  compounds,  are  also  produced  by 
the  metals  adjacent  to  the  metalloids  : 

II  III  IV  III  II  I 

MgfCIb),  Al(CIl3)3  Si{CH3h  P(CH3)3  S(CIl3),  CICIL. 


248 


INORGANIC  CHEMISTRY. 


Their  Stability  gradually  dimill  islics  with  the  increasing  basic  nature  of 
the  metals;  hence,  in  the  three  large  jieriods,  this  power  extends  only  to 
Zn,  Cd,  and  Hg. 

In  consequence  of  the  opposite  character  of  the  two  ends  of  the  jieriods, 
there  are  in  the  table  representing  the  double  ])eriodicily  of  the  great 
periods  (pp.  244  and  246)  two  sub-groups  each,  with  the  seven  vertical 
groups ; on  the  left  with  the  more  positive,  basic,  and  on  the  right  with 
the  more  negative,  metalloidal  elements.  Thus  in  group  VI,  in  addition 
to  O and  S (belonging  to  the  small  periods)  stands  the  more  basic  sub- 
group Cr,  Mo,  W,  and  the  metalloids  Se  and  Te  ; in  group  II  stand  the 
strong  basic  metals  Ca,  Sr,  Ba,  and  the  less  basic  heavy  metals  Zn,  Cd, 
Hg.  The  elements  of  group  VIll  form  the  gradual  transition  from  the 
latter  to  the  former. 


Periodicity  of  Chemical  Valence. — Groui)  I of  the  table  com- 
prises the  univalent  metals,  group  II  the  bivalent.  In  group  III  are  the 
trivalent  metalloid,  boron,  and  the  tri valent  metals  Al,  Sc,  Y,  and  Ga, 
In,  Tl.  In  the  quadrivalent  carbon  group  the  valence  arrives  at  its  maxi- 
mum ; trom  here  it  gradually  decreases  with  increasing  atomic  weight ; 
the  nitrogen  group  is  trivalent,  the  oxygen  group  is  bivalent,  that  of  the 
halogens  univalent.  This  valence  is  derived  from  the  compounds  with 
hydrogen  and  hydrocarbons  (compare  p.  247),  or  where  such  do  not 
exist,  as  in  the  case  of  boron  and  many  metals,  from  the  halogen  com- 
pounds : 


IV 

III 

II 

I 

CH, 

NH3 

UH2 

FIH 

I 

II 

III 

IV 

III 

II 

I 

LiCl 

BeCq 

Bcq 

ccq 

NCI3 

0CI3 

Fb 

NaCl 

MgCb 

AICI3 

Sicq 

PCI3 

seq 

cq. 

The  elements  of  the  first  four  groups  are  not  capable  of  yielding  higher 
compounds  with  the  halogens.*  On  the  other  hand,  as  we  have  seen, 
the  higher  analogues  of  nitrogen  and  other  metalloids  can  unite  with  a 
larger  number  of  halogen  atoms  (see  pp.  169,  1 70).  The  higher  valence  of 
these  elements  is  more  manifest  in  the  more  stable  oxygen  compounds. 
On  bringing  together  the  highest  oxides  of  the  seven  groups  capable  of 
forming  salts  (salt-building  oxides),  we  get  this  series: 

I II  III  IV  V VI  VII 

LijO  BeO  COj  SO3  (ip/). 

The  elements  of  the  first  four  groups  in  their  oxygen  compounds 
exhibit,  consequently,  the  same  valence  as  in  the  compounds  with  hydro- 
gen (or  hydrocarbon  radicals)  and  the  halogens;  in  the  last  three  series, 
however,  there  is  noticed  a constant  increase  of  valence  for  oxygen. f 


* Recently  lialide  derivatives  of  c;esium  and  rubidium  of  the  formulas  CsMj  and 
CsMr,  liave  been  prej)ared  (see  Oesium). 

f 'I'lie  comix)und  sulphur  heptoxide  (jn  188),  evidently  does  not  fit  in  here.  It 

v(uy  probably  belongs  to  the  (rue  peroxides  with  the  chain  -()-()-,  so  that  its  compo- 
sition does  not  give  a definite  answer  as  to  the  valence  of  sulphur. 


PERIODIC  SYSTEM  OF  THE  ELEMENTS. 


249 


Besides  these  highest  oxides,  remarkable  for  their  greater  stability,  the 
elements  of  the  last  three  groups  afford  lower  oxides,  returning  in  this 
manner  to  the  hydrogen  valence : 


III 

IV 

V 

P303 

S02 

bOs 

II 

III 

SCb 

(CIA) 

I 

CI2O 

PH3 

SHj 

cm. 

The  hydroxyl  compounds  of  the  elements  of  the  seven  groups  are  anal- 
ogous to  the  oxides  in  constitution.  They  afford  the  following  series, 
expressing  the  maximum  valence  (compare  p.  169)  : 

I II  III  IV  V VI  VII 

Na(OH)  Mg(OH)2  Al(OH)3  Si(OHh  P(OH)5  S(OH)e  0(011)^. 

The  hydroxyl  compounds  of  the  elements  of  the  first  four  groups  exist  in 
free  condition,  excepting  that  of  carbon,  C(OH)^,  which  is  only  repre- 
sented in  its  derivatives.  . The  strong  basic  character  of  the  hydroxides 
of  group  I (NaOH)  diminishes,  step  by  step,  in  the  succeeding  groups, 
down  to  the  weak  acid  hydrate,  Si(OH)^.  The  hydrates  of  the  last  three 
groups  are  of  acid  nature,  and  mostly  unstable  or  not  known.  By  the 
elimination  of  one,  two  and  three  molecules  of  water  they  yield  the 
ordinary  acids: 

V VI  VII 

P0(0H)3  S02(0H)2  C103(0H). 

Phosphoric  acid.  Sulphuric  acid.  Perchloric  acid. 

The  non-saturated  hydroxides  behave  in  the  same  way : 

III  IV  V 

?(0H)3  S(0H),  C1(0H)3 

III 

C1(0H)3 

I 

. . Cl(OH). 

Sulphurous  acid,  S0(0H)2,  is  derived  from  the  hydrate,  S(OH)j 
chloric  acid,  ClOg.OH,  from  the  hydrate,  Cl(OH)5;  and  chlorous  acid, 
CIO. OH,  from  the  hydrate,  Cl(OH)3.  The  hydrates,  P(0H)3  and 
ClOH,  are  very  unstable,  and  the  first  appears  to  pass  readily  into 
HP0(0H)2  (compare  p.  215). 

It  has  been  already  shown  in  the  case  of  periodic,  sulphuric,  and  nitric 
acids,  how  the  so-called  hydrates  with  water  of  crystallization  (regarded 
as  molecular  compounds)  are  explained  by  the  acceptance  of  the  exist- 
ence of  such  hydroxyl  derivatives.  The  same  may  be  done  for  many  salts 
with  water  of  crystallization. 

Thus,  we  see,  and  in  the  following  pages  will  find  it  more  extensively 
developed,  that  the  relations  of  valence  of  the  elements  have  their  com- 
plete expression  in  the  periodic  system,  are  regulated  by  it,  and  hence 
we  must  conclude  that,  in  fact,  the  valence  is  not  only  a projierty  at- 
taching to  the  elements  per  se,  but  is  influenced  also  by  the  nature  of  the 


250 


INORGANIC  CHEMISTRY. 


combining  elements;  the  hydrogen  valence  is  constant,  the  valence  of 
oxygen  and  the  halogens,  on  the  contrary,  varies  according  to  definite 
rules.  Valence,  therefore,  is  a relative  function  of  the  elements  (j).  169). 

The  periodic  system,  the  natural  division  of  the  elements,  apjiarently 
emphasizes  the  fact  that  the  atoms  of  the  chemical  elements  are  aggrega- 
tions or  condensations  of  one  and  the  same  iirimordial  substance — the 
unity  of  matter,  which  corresponds  to  the  unity  of  force  already  known 
tons.  It  is  only  by  assuming  a jirimordial  substance  that  the  periodic 
dependence  of  the  properties  of  the  chemical  elements  iiiion  the  magni- 
tude of  the  atomic  weights  can  be  comjirehended. 

It  was  suggested  that  hydrogen  was  this  original  substance,  because  it 
appeared  as  if  all  atomic  weights  had  to  be  expressed  in  whole  numbers 
— that  is,  multiples  of  the  atomic  weight  of  hydrogen,  taken  as  unit 
(Front  1815,  Meinecke  1818).  More  accurate  determinations,  made 
with  exceeding  care  by  Stas,  demonstrated  this  not  to  be  true  in  all 
cases  [J.  S.  Stas,  Untersuchungen  iiber  die  Gesetze  der  chemischen  Pro- 
portionen,  iiber  die  Atomgewichte  und  ihre  gegenseitigen  Verhaltnisse, 
Deutsch  von  Aronstein,  Leij)zig,  1867].  The  inquiry  was  then  made  as 
to  whether  the  variations  of  the  atomic  weights  from  whole  numbers 
could  not  be  occasioned  by  definite  amounts  of  ])onderable  ether  enter- 
ing or  escaping  in  the  chemical  transposition  of  bodies.  The  experi- 
ments of  H.  Landolt  on  this  point  gave  a negative  reply,  thus  cutting 
off  the  last  avenue  which  remained  open  to  the  hypothesis  of  Prout  and 
Meinecke  [Ber.  26  (1893),  ii,  1820]. 

It  must  be  borne  in  mind,  however,  that  it  is  not  only  the  atomic 
weight  which  influences  the  properties  of  an  element,  but  that  the  molec- 
ular weight  and  the  energy  of  the  molecule  exercise  great  influence  in 
this  direction.  The  occurrence  of  allotropic  modifications  of  an  element 
can  only  be  explained  on  this  assumption.  These,  in  their  general 
deportment,  often  differ  more  from  one  another  than  two  different  but 
related  elements.  Phosphorus  in  its  varieties  is  a most  striking  example. 
Similar  relations  have  been  noticed  with  carbon — soot,  graphite,  dia- 
mond, and  with  oxygen — ordinary  oxygen  and  ozone.  It  is  only  the 
power  of  uniting  a definite  number  of  different  atoms  which  remains  in 
all  varieties  of  one  element.  Thus,  red  phosphorus  on  oxidation  yields 
the  same  oxides  and  acids  as  the  yellow,  and  by  burning  the  diamond  the 
same  carbonic  acid  is  produced  as  with  soot  and  graphite.  The  valence 
which  one  element  shows  with  reference  to  other  elements,  is,  there- 
fore, primarily  a function  of  the  atomic  weight.  Its  physical  properties, 
on  the  other  hand,  are  also  dependent  upon  the  molecular  weight  and  the 
energy  of  the  molecule. 

The  la7v  ttnderlyifig  the  periodic  system  will  only  be  recognized  perfectly 
and  in  its  eiitirety  ivhcn  it  becomes  possible  to  deduce  the  properties  of  the 
different  varieties  of  a )i  element  fro7n  the  atomic  7veight  of  the  latter. 

Again,  these  allotrojiic  modifications  indicate  that  bodies  can  be  formed 
by  a different  arrangement  of  similar  atoms,  and  it  is  only  by  keen  scien- 
tific observation  that  varieties  of  one  and  the  same  substance  can  be 
detected,  d’he  elements  could  similarly  consist  of  varieties  or  allotropic 
forms  of  some  one  ])rim()rdial  form  of  matter. 


THE  METALS. 


Although  there  is  no  sharp  line  of  demarcation  between  metals  and 
non-metals,  yet  these  two  classes  of  bodies  are  characteristically  distinctive 
in  their  entire  deportment,  as  may  be  plainly  seen  in  the  periodic  system 
of  elements.  In  physical  respects  the  character  of  metals  is  determined 
by  their  external  appearance  and  by  their  ability  to  conduct  heat  and 
electricity;  chemically,  it  shows  itself  chiefly  in  the  basicity  of  the  oxygen 
compounds ; yet  we  see  that  with  the  increase  of  the  number  of  the  oxygen 
atoms,  the  basic  character  gradually  diminishes  and  becomes  acidic. 
Gaseous  compounds  of  the  metals  with  hydrogen  are  not  known. 


PHYSICAL  PROPERTIES  OF  THE  METALS. 

At  ordinary  temperatures  all  the  metals  excepting  mercury  are  solid, 
slightly  volatile  bodies.  They  are  opaque,  and  only  a few,  like  gold, 
permit  the  passage  of  light  to  a limited  extent  when  beaten  into  thin 
leaflets.  In  compact  mass  they  exhibit  metallic  luster  and  mostly  possess 
a whitish-gray  color ; gold  and  copper  are,  however,  brilliantly  colored. 
In  powder  form  almost  all  the  metals  are  black. 

Most  of  them  crystallize  in  the  forms  of  the  regular  system ; only  a 
few,  showing  a metalloidal  character,  are  not  regular.  Thus  antimony 
and  bismuth  crystallize  in  the  hexagonal  system,  and  tin  is  quadratic. 

The  specific  gravities  of  the  metals  vary  greatly,  from  0.59  to  22.5,  as 
seen  from  the  following  arrangement : 


Ij'thium, 

0-59 

Germanium, 

5*47 

Ixad, 

11-37 

Pota.ssium, 

0.87 

Zinc, 

7-1 

Palladium, 

II-5 

Sodium, 

Rubidium, 

0.97 

Tin, 

Iron, 

7-3 

Thallium, 

II. 8 

152 

7.8 

Mercury, 

13-55 

Calcium, 

1-57 

Cobalt, 

8.5 

Gold, 

19-3 

Magne.sium, 

1.74 

Copper, 

8.9 

Platinum, 

21.5 

Aluminium, 

2.6 

Bismuth, 

9.8 

Iridium, 

Osmium, 

22.4 

Barium, 

3-75 

Silver, 

10.5 

22.5 

In  general  the  specific  gravities  of  the  metals,  and  also  those  of  the 
metalloids,  increase  with  the  atomic  weights;  they  stand  more  especially 
in  sharp  periodic  dependence  with  reference  to  the  latter.  The  first 
members  of  all  periods  possess  low  specific  gravities;  the  latter  grow 


* Liquid  ; solid  = 14. 2. 
251 


252 


INORGANIC  CHEMISTRY. 


gradually  until  the  middle  of  the  period,  when  the  maximum  is  attained, 
and  then  they  again  decrease  (p.  245).  These  relations  show  themselves 
more  fully  if,  instead  of  the  specific  gravity,  we  comi)are  the  siiecific 
volumes  or  atomic  volumes;  i.  e.,  the  cpiotients  from  the  atomic  weights 
(A)  and  specific  gravities  (d) : 

^ = specific  volume. 

These  quotients  express  the  relative  value  of  the  atoms  (in  solid  or 
liquid  state),  'hhus  the  atomic  volume  of  lithium  = 11.9,  that 
of  potassium  = 45-o^  i-  potassium  atom  occiqn'es  a sjiace 

3.8  times  larger  than  that  of  the  lithium  atom.  The  periodic  alterations 
of  the  atomic  values  are  in  opposition  to  those  of  the  specific  gravities,  as 
the  former  are  obtained  by  the  division  of  the  atomic  weights  by  the 
specific  gravities.  Therefore,  the  atomic  volumes  decrease  gradually, 
commencing  with  the  first  members  of  the  periods,  attain  a minimum  in 
the  middle  of  the  periods,  and  then  increase  again  up  to  the  last  mem- 
bers. On  the  other  hand,  we  find  that  with  the  homologous  elements 
(the  vertical  series)  an  increase  in  the  atomic  volumes  almost  invariably 
occurs  as  the  atomic  weights  increase. 

The  metals  whose  specific  gravities  are  less  than  5 are  termed  light 
7netals,  the  rest  heavy  7netals.  The  naturally  occurring  compounds  of  the 
heavy  metals  are  termed  ores. 


Most  metals  are  very  77talleable  and  tenacious,  hence  can  be  beaten  into 
thin  plates  and  leaves,  and  drawn  out  into  wires ; gold  and  silver  are  the 
most  malleable.  A few,  like  bismuth  and  tin,  possess  a metalloidal  char- 
acter, are  brittle,  and  may  be  pulverized  like  antimony  and  arsenic. 

Heat  will  fuse  all  metals,  although  some  require  high  temperatures. 
The  fusing  points  of  the  most  important  of  them  are  the  following : 


Mercury,  . . . , 

. . . . —39° 

Aluminium.  . . . 

• . . 650° 

Rubidium,  . . . , 

. . . .+38° 

Germanium,  . . . 

. . . 900° 

Potassium,  . . . , 

. . . . 62° 

Silver, 

. . . 960° 

Sodium,  , . , , 

. . . . 96° 

Gold, 

. . . 1060° 

Tin, 

. . . . 230° 

Copper, 

. . . 1090° 

Bismuth,  . . . , 

. . . . 265° 

Cast  iron,  .... 

. . . 1150° 

Cadmium,  . . . , 

. . . . 320° 

Wrought  iron,  . . 

. . . 1600° 

Lead,  

. . . . 330° 

Platinum,  .... 

. . . 1775° 

Zinc, 

...  420° 

Iridium,  .... 

. . . 1950° 

Ruthenium,  osmium  and  also  chromium,  molybdenum,  uranium,  tung- 
sten and  vanadium  melt  with  greater  difficulty  than  platinum. 

A greater  volatility  also  corresponds  to  the  greater  fusibility.  Mer- 
cury boils  at  357°;  potassium  and  sodium  at  about  700°;  cadmium  at 
770°;  zinc  about  950°,  and  the  difficultly  fusible  metals  may  be  also 
volatilized  in  the  electric  furnace. 

Metallic  borides,  silicides  and  carbides  (p.  163)  withstand  temperatures 
which  volatilize  i)latinum,  calcium  and  carbon  (Moissan). 


SPECIFIC  HEAT — ATOMIC  HEAT.  253 

All  these  physical  properties  bear  a periodic  dependence  to  the  atomic 
weights,  as  will  be  more  plainly  indicated  in  the  indvidual  groups. 

Until  lately  the  highest  temperature  attainable  for  scientific  and  technical  uses  was 
about  2000°.  The  idea  of  using  the  electric  arc  to  extend  the  range  of  temperature  is 
not  new,  but  the  actual  achievement  was  first  accomplished  by  the  French  chemist,  H. 
Moissan,  by  the  construction  of  the  electric  furnace.  The  principle  involved  in  the  latter 
is  to  expose  a substance  to  the  heat  of  the  electric  arc  in  a space  as  small  as  possible  and 
surrounded  by  fire-resisting  material.  The  action  is  then  due  to  heat  and  not  to  elec- 
tricity. Moissan  in  this  way  obtained  a temperature  of  3500°  C.  The  furnace  model 
first  used  by  him  in  1892  is 
sketched  in  Fig.  68.  It  con- 
sists of  two  blocks  of  unslaked 
lime.  The  lower  contains  a 
groove  for  the  carbon  elec- 
trodes E,  E,  and  about  the 
middle  is  a depression  in 
which  the  material  is  placed 
and  upon  which  the  heat  of 
the  arc  is  to  act.  The  material 
can  also  be  exposed  in  a small 
carbon  crucible.  The  upper 
portion  is  slightly  arched  above  the  arc.  As  the  lime  soon  melts  on  the  surface  and 
becomes  soft  the  arch  acts  as  a heat  reflector.  For  the  various  forms  of  electric  furnaces 
consult  H.  Moissan,  Der  elekt.  Ofen,  1897  ; also  Liebetanz.  Calciumkarbid-  und  Acety- 
lentechnik,  pp.  38-79.  Compare  Z.  f.  anorg.  Ch.  22  (1896),  293. 

SPECIFIC  HEAT.  ATOMIC  HEAT. 

Of  all  physical  properties  of  the  elements,  from  a chemical  standpoint, 
their  heat  capacity  is  the  most  important,  as  it  can  serve  for  the  determi- 
nation of  the  atomic  weights.  The  specific  heat  of  a substance  is  the 
quantity  of  heat  which  increases  its  mass  unit  i°  in  temperature.  Heat 
quantities  are  measured  in  either  large  or  small  calories  (p.  66),  and 
according  as  one  or  the  other  is  chosen,  the  expression  for  the  mass  unit 
will  be  I kilogram  or  i gram.  As  water,  of  all  substances,  possesses  the 
greatest  specific  heat,  and  as  this  also  serves  as  unit;  therefore,  for  all 
other  bodies  the  specific  heat  represents  the  fraction  of  a calorie  which 
raises  their  mass  unit  i°. 

The  atomic  heat  of  an  element  is  the  product  of  its  specific  heat  (H) 
and  atomic  weight  (A).  Dulong  and  Petit  (i8i8)  discovered  the  remark- 
able fact  that  the  solid  elements  possess  approximately  the  same  atomic  heat. 
It  is  about  6.4  : 

A.  H = 6.4. 

For  silver  A = 107.93  H = 0.057  ; A . H = 6.  i ; for  lead  A = 
206.9,  H = o.o3i,A.H  = 6 4;  for  potassium  A = 39.15,  H = 0.166, 
A.  H=  6.5 — /.  e.,  107.93  gi'ams  of  silver,  206.9  grams  of  lead,  and  39.15 
grams  of  potassium  are  heated  by  approximately  6.4  small  calories 
through  1°  of  temperature,  or  107.93,  206.9  and  39.15  kilograms  are 
raised  1°  by  6.5  large  calories.  The  equation  A . H = 6. 4 also  indicates 
that  the  specific  heats  of  two  elements  are  to  each  other  approximately  as 
their  atomic  weights — the  greater  the  atomic  weight,  the  smaller  the 
specific  heat. 

It  is  only  in  the  case  of  a few  of  the  elements  that  the  atomic  heat 


254 


INORGANIC  CHEMISTRY. 


varies  widely  from  the  mean  6.4,  e.  g.,  sulphur  is  5.7,  phosphorus  5.9, 
silicon  4.6,  germanium  5.6,  while  carbon  is  2-2.8,  boron  2.6  and  beryl- 
lium 3.7.  These  elements  have  low  atomic  weight  and  are  either  metal- 
loids or  resemble  metalloids. 

These  variations  from  the  mean  are  in  part  explained  by  the  change  in 
s])ecific  heat  with  the  temperature.  It  was  known  before  that  these  show  a 
slight  increase  with  the  temperature,  but  it  is  only  recently  (1875)  that 
Weber  has  proved  that  the  increase  is  very  considerable  for  the  elements 
C,  B and  Si,  which,  at  medium  temperatures,  possess  a remarkably  low 
atomic  heat ; that,  beyond  a definite  temperature,  the  atomic  heat  becomes 
tolerably  constant,  and  then  almost  agrees  with  the  law  of  Dulong  and 
Petit.  According  to  Nilson,  beryllium  shows  a similar  deportment  : 


H 

A 

H X A 

Diamond,  graphite,  above  900°, 

0.459 

12.00 

5-5 

Boron,  above  600°,  

0-5 

I I 

5-5 

Silicon,  above  200°, 

0.204 

28.4 

5 8 

Beryllium,  above  257°, 

0.58 

9.1 

5 3 

From  this  close  agreement  of  the  found  atomic  heat  of  the  metals  with 
the  mean,  it  follows,  without  doubt,  that  there  does  occur  a regularity, 
and  we  must  conclude  that  the  slight  variations,  apart  from  the  inaccuracy 
of  the  observations  and  the  impurities  of  the  substances  used,  are  influ- 
enced by  physical  causes  of  an  unknown  nature.  Hence,  the  specific  heat 
may  serve  for  the  derivation  of  the  atomic  weight  of  the  elements;  ^/le 
atomic  weight  is  equal  to  the  eo?istafit  6.4  divided  by  the  found  specific  heat : 


The  atomic  weights  derived  from  the  specific  heat — the  so-called  ther- 
mal atomic  weights — agree  in  almost  all  instances  with  those  obtained 
from  the  vapor  density  of  the  free  elements  or  their  volatile  compounds. 
Where  no  volatile  compounds  of  an  element  are  known,  the  specific  heat 
is  the  only  certain  means  of  fixing  the  actual  atomic  weight.  The  equiv- 

I 

alent  weight,  38  (InCl),  of  indium  was  fixed  with  great  accuracy  by 
analysis ; it  was,  however,  unknown  whether  the  atomic  weight  was 
double  or  triple  that  quantity.  The  specific  heat  of  indium  was 
found  to  be  0.0569,  from  which  the  atomic  weight  would  be 
= 112,  a number  closely  ap})roaching  the  trebled  equivalent  weight  of 
indium,  114  (=  38  X 3)-  From  this  it  follows  that  the  true  atomic 
weight  of  indium  is  114  and  that  indium  is  trivalent  (InClg). 

In  their  solid  com])ounds  the  elements  retain  the  specific  heat  ; hence  the  molecular 
heat  is  nearly  e(|ual  to  the  sum  of  the  atomic  heats  of  the  elements  constituting  the 
molecules — law  of  h'.  Neumann  and  II.  Ko])]).  I lence  the  atomic  heat  of  elements  not 
known  in  solid  condition  may  be  derived  from  the  molecular  heat  of  their  compounds. 
In  this  manner  the  following  atomic  heats  are  found  : For  nitrogen,  5-0  j for  chlorine, 
C ; for  oxygen,  4 ; for  (luorine,  5 ; for  hydrogen,  2.3  ; for  carbon,  1. 8. 


CHEMICAL  PROPERTIES  OF  THE  METALS.  255 

In  the  free  state  the  gaseous  elements  usually  have  a slighter  atomic  heat,  as  seen  from 
the  following  table  : 


A 

H* 

A X H 

Oxygen, 

16 

0. 156 

2.5 

Hydrogen, 

1 .01 

2.405 

2.4 

Nitrogen, 

14.04 

0.172 

2.4 

Chlorine, 

3545 

0.093 

3 3 

ISOMORPHISM. 

As  irrdicated  in  the  preceding  pages,  the  atomic  weights  of  the  elements  may  be 
derived  directly  from  the  specific  heat  of  solids,  while  from  the  gas  density  of  the  volatile 
compounds  we  get  the  molecular  weights,  and  from  the  latter,  indirectly,  ascertain  the 
atomic  weights  (compare  p.  80).  A third,  although  less  general  and  certain,  means  of 
determining  the  atomic  weight  is  afforded  by  isomorphism.  By  this  is  understood  the  phe- 
nomenon observed  by  Mitscherlich  (1819).  Characteristics  of  the  isomorphism  of  two 
bodies  are  : 

(l)  Similarity  in  crystalline  form,  whereby  the  symmetry  properties  agree  perfectly  and 
the  geometrical  constants  nearly  coincide  (see  pp.  31,  38)  ; (2)  the  power  of  forming 
within  certain  definite  limits  mixed  crystals  in  any  desired  quantity  by  weight  without 
altering  the  crystal  form;  (3)  the  power  of  overgrowth,  i.  e.,  crystals  of  the  one  sub- 
stance being  able  to  continue  their  growth  in  supeisaturated  solutions  of  the  other  com- 
pound. From  Mitscherlich’ s discovery  we  can  conclude  from  the  isomorphism  of  two 
compounds  that  they  are  similarly  constituted  and  have  an  equal  number  of  atoms  in  the 
molecule.  Accordingly  the  quantities  of  the  elements,  occurring  in  isomorphous  mixtures, 
would  be  to  each  other  as  their  atomic  weights,  and  this  would  make  it  possible  to  estab- 
lish an  incorrect  atomic  value.  For  example,  the  metals  calcium,  strontium  and  barium 
do  not  afford  volatile  derivatives.  Their  atomic  weights  could  not  be  deduced  from  their 
thermal  capacity,  and  it  was  the  isomorphism  of  many  of  their  compounds  with  those  of 
magnesium  that  determined  the  same  ; the  quantities  of  these  elements,  replacing  24.36 
parts  by  weight  of  magnesium  (l  atom),  were  accepted  as  the  true  atomic  weights. 

In  the  present  state  of  chemistry  we  attach  but  secondary  importance  to  isomorphism  as 
a method  of  determining  atomic  weights  because  it  is  frequently  very  difficult  to  demon- 
strate that  isomorphism  is  actually  present,  for  the  coincidence  in  geometrical  constants 
requires  to  be  only  approximate,  the  ability  to  yield  mixed  crystals  occurs  in  every  imag- 
inable degree  and  bodies  possessing  not  the  slightest  chemical  similarity  do  show  over- 
growths. It  is  indeed  true  that  bodies  similar  chemically  often  exhibit  like  crystalline 
form,  and  the  power  of  forming  mixed  crystals  also  increases  the  greater  the  similarity  in 
chemical  properties. 


CHEMICAL  PROPERTIES  OF  THE  METALS. 

Alloys. — Many  metals  may  be  so  mixed  by  fusing  them  together  that  a 
solid,  homogeneous  mass  will  result  upon  cooling ; the  constituents  cannot 
be  separated  then  by  mechanical  means.  Such  a mixture  is  based  upon  a 
molecular  interpenetration  ; it  cannot  be  distinguished  from  a mere 
mechanical  mixture  and  it  bears  the  name  alloy  \legere,  put  together). 


* By  constant  volume. 


256 


INORGANIC  CHEMISTRY. 


A great  deal  of  heat  is  frequently  developed  just  as  in  a chemical  union 
by  fusing  metals  or  by  bringing  together  those  already  melted.  Yet  alloys 
in  general  cannot  be  considered  chemical  compounds;  their  composition, 
even  in  the  case  of  those  which  crystallize,  is  rarely  according  to  the 
atomic  ratio.  They  should  rather  be  regarded  as  solid  solutions  of  one 
metal  in  another  or  of  metallic  comi)ounds  in  a metal,  if,  following 
Ostwald,  a homogeneous  mixture  of  various  substances  (showing  no  vari- 
ability) is  defined,  without  reference  to  its  state  of  aggregation,  as  a 
solution. 

Many  molten  metals  may  be  mixed  just  like  water  and  alcohol  in  every 
l)roportion  with  one  another,  e.  g.,  tin  and  lead  ; others  in  a less  degree, 
similarly  to  ether  and  water,  e.  g.,  zinc  and  lead,  d'he  second  class  of 
alloys  occurs  more  frequently.  Usually  metals  which  are  similar  chemic- 
ally alloy  more  readily  with  one  another  than  those  which  exhibit  marked 
differences  in  their  chemical  behavior. 

Many  alloys  can  be  obtained  crystallized,  yet  in  actual  practice  crys- 
tallization is  avoided  because  it  diminishes  the  solidity,  ductility,  etc. 

As  a rule,  the  density  of  an  alloy  cannot  be  calculated  from  its  constit- 
uents, but  this  is  possible  with  a mere  mechanical  mixture.  Many  alloys 
exhibit  a very  definite  condensation  (contraction),  e.  g.,  silver-gold,  cop- 
per-tin ; others  a distinct  expansion:  copper-silver,  antimony-tin,  tin- 
lead,  while  others  do  not  manifest  either  of  these  properties:  copper-gold, 
antimony-bismuth. 

The  color  of  alloys  does  not  correspond  to  the  quantity  of  their  con- 
stituents, for  copper  alloys,  with  30  per  cent,  of  tin,  are  white  in  color,  and 
pure  yellow  with  the  same  amount  of  zinc.  The  bright  color  disappears 
with  60  per  cent,  of  zinc.  Aluminium  changes  the  color  of  copper 
readily;  the  same  is  true  of  nickel,  as  may  be  seen  in  the  coin,  contain- 
ing about  25  per  cent,  of  nickel  and  75  per  cent,  of  co])per.  Gold  possesses 
slight  coloring  power.  Gold  and  silver  alloys  having  30  per  cent,  of 
gold  possess  a pure  silver-white  color. 

The  melting  point  of  an  alloy  is  generally  lower  than  would  be  expected 
from  the  melting  points  of  its  constituents.  An  alloy  of  8 parts  of  lead,  15 
parts  of  bismuth,  4 parts  of  tin,  and  3 parts  of  cadmium  melts  at  68”, 
although  each  of  these  metals  melts  above  200”.  Potassium  (m.  p. 
62°)  and  sodium  (m.  p.  96”)  unite  in  the  ratio  of  their  atomic  weights 
to  an  alloy  which  is  liquid,  like  mercury,  at  the  ordinary  temperature. 

Quite  frequently  molten  alloys  manifest  the  tendency  on  slow  cooling 
to  separate  into  several  compounds  which  differ  in  composition,  fusibility, 
etc.,  from  one  another.  Under  these  conditions  pure  metals  also  sepa- 
rate from  many  alloys.  This  process,  very  important  for  the  preparation 
of  metals,  is  designated  segregation  (see  Silver). 

Alloys  often  resist  acids  much  better  than  single  metals;  however,  they 
frequently  dissolve  in  acids  even  when  the  one  constituent  is  insoluble  in 
them.  Thus  an  alloy  of  56. 5 i)cr  cent,  of  copper  and  43.5  ])er  cent,  of  zinc 
is  scarcely  attacked  by  nitric  acid,  and  in  gold-silver  alloys,  where  there 
is  a prejiondcrance  of  gold,  the  silver  is  not  dissolved  out  by  nitric  acid. 
If  the  alloy  contains  3 ])arts  of  silver  to  one  part  of  gold  the  nitric  acid 
will  dissolve  the  silver  (cpiartation). 


CHEMICAL  PROPERTIES  OF  THE  METALS. 


257 


Mercury  is  able  to  dissolve  almost  all  metals,  forming  alloys  known  as 
aftialgains  (from  imlayiia,  a soft  plaster),  which  are  generally  crystallizable. 
In  chemical  respects  hydrogen  is  a metal,  and  very  many  of  the  metals 
combine  with  it.  Palladium,  potassium  and  sodium  yield  the  com- 
pounds K^H2  and  Na^H.^,  which  deport  themselves  as  alloys.  That 
antimony  yields  a gaseous  product  (SbHg),  is  due  to  its  pronounced 
metalloidal  character.  The  ability  of  individual  metals  of  the  platinum 
and  iron  groups  to  permit  the  passage  of  hydrogen  at  a red  heat  depends, 
probably,  upon  a chemical  attraction;  hydrogen  first  dissolves  and  is 
then  evaporated  again. 

Metallic  Carbides. — These  are  produced  by  dissolving  carbon  in 
the  molten  metal  or  by  reducing  metallic  oxides  with  carbon.  The  elec- 
tric furnace  with  its  high  temperatures  has  been  applied  in  this  direction. 
The  researches  of  Moissan  and  the  technical  application  of  calcium  car- 
bide and  silicon  carbide  (carborundum)  have  given  great  theoretical  and 
practical  importance  to  the  carbides  in  general  (pp.  154,  163).  Molten 
gold,  bismuth,  lead  and  tin  do  not  dissolve  carbon,  and  copper  and  silver 
very  little.  The  melted  platinum  metals  take  it  up  in  large  amounts,  and 
upon  cooling  it  separates  as  graphite.  Many  other  metals,  under  like 
conditions,  form  carbides  which  crystallize  beautifully.  Carbides  may 
be  arranged  in  two  series: 

1.  Those  decomposed  by  water  at  the  ordinary  temperature.  Here 
belong  the  carbides  of  potassium,  lithium,  Li2C2,  calcium,  barium,  stron- 
tium (Me"C2),  aluminium,  Al^Cg,  manganese,  MnC,  and  beryllium. 

2.  Those  which  are  not  affected  by  water  : molybdenum  carbide,  CMOg, 
chromium  carbide,  CrgCg,  etc.  The  carbides  decomposable  by  water 
yield  hydrogen,  methane,  acetylene,  and  also  liquid  and  solid  hydrocar- 
bons. Moissan  thinks  that  this  behavior  affords  hints  as  to  the  origin 
of  petroleum  (see  Calcium-  and  Aluminium  Carbide,  also  p.  252). 


Halogen  Compounds. — The  metals  unite  directly  with  the  halogens 
to  form  compounds,  which  are  not  decomposed  by  water  at  ordinary  tem- 
peratures, and,  in  general,  are  very  stable ; on  the  other  hand,  the  halogen 
compounds  of  the  metalloids  (excepting  those  of  carbon)  are  easily 
decomposed  by  water.  These  compounds  are  also  produced  by  the  action 
of  the  haloid  acids  upon  the  free  metals,  their  oxides,  hydroxides,  and 
carbonates,  whereby  they  plainly  characterize  themselves  as  salts  of  the 
haloid  acids.  A third  procedure  for  the  formation  of  chlorides  and  bro- 
mides, essentially  analogous  to  the  first,  is  based  upon  the  simultaneous 
action  of  carbon  and  chlorine,  or  bromine  upon  the  oxides  (see  Aluminium 
Chloride  and  Silicon  Chloride). 

The  following  types  of  halogen  derivatives  exist  and  show  the  different 
valences  of  the  metals : 

I II  III  IV  V VI 

KCl  ZnClj  InClg  SnCl,  TaClg  WClg. 


22 


258 


INORGANIC  CHEMISTRY. 


OXIDES  AND  HYDROXIDES— HYDRATES. 

I'he  affinity  of  the  metals  for  oxygen  varies.  Some  of  them  oxidize  in 
moist  air  and  decompose  water,  even  at  ordinary  temperatures.  Such  are 
the  so-called  alkalies  and  alkaline  earths  (the  potassium  and  calcium 
groups).  Their  oxides  dissolve  readily  in  water  and  form  strong  basic 
hydroxides  or  hydrates,  KOH,  Ca(OH)2,  which  are  usually  not  decom- 
])Osed  by  ignition. 

Other  metals  (the  so-called  heavy  metals)  oxidize  and  decompose  water 
only  at  higher  temperatures;  their  oxides  are  insoluble  in  water,  and 
generally  afford  no  hydroxides,  as  the  latter  ui)on  heating  readily  decom- 
])Ose  into  oxides  (anhydrides)  and  water: 

Zn(OII)2  = ZnO  + H^O. 

They  are  of  a less  basic  nature,  and  their  soluble  salts  usually  exhibit 
acid  reaction.  Finally,  some  metals,  as  gold  and  platinum  (the  noble 
metals),  are  incapable  of  combining  directly  with  oxygen.  Their  oxides, 
obtained  in  another  way,  decompose  readily  under  the  influence  of  heat 
into  metal  and  oxygen.  The  universal  method  for  the  preparation  of 
insoluble  oxides  and  hydroxides  of  the  heavy  metals  depends  upon  the 
precipitation  of  the  solutions  of  their  salts  by  alkaline  bases: 

HgCl.,  + 2KOII  = 2KCI  -h  HgO  + HjO 

II  II 

CuSO*  + 2KOH  = K2SO4  + Cu(OH)j. 


The  different  valences  of  the  metals  are  most  clearly  seen  in  their  oxygen  derivatives, 
which  form  salts.  We  have  the  following  eight  forms  or  types  of  the  highest  salt-pro- 
ducing oxides  (see  p.  244),  corresponding  to  the  eight  groups  of  the  periodic  system  of 
the  elements  : 

I II  III  IV  V VI  VII  VIII 

K2O  MgO  AI2O3  Sn02  Bi205  CrOg  Mn207  OSO4. 

The  hydroxides  of  the  first  two  forms  possess  a strong  basic  character  and  only  yield 
salts  with  acids.  In  the  hydroxides  of  the  two  succeeding  forms  there  is  shown  an  acid- 
like character  together  with  the  predominating  basic  character.  Hence  they  dissolve  in 
alkalies  and  form  salt-like  derivatives  with  bases,  in  which  hydrogen  is  replaced  by  metals, 
e.  g.,  Al(ONa)3.  These  higher  (normal)  hydrates  are  not  very  stable,  give  up  water  and 
})ass  into  meta-hydrates,  which  retain  the  acid  character.  Thus,  from  Al(OH)3is  derived 
AlO.OH,  which  yields  salt-like  compounds,  e.  g.,  AlO . ONa ; from  Sn(OH)4  are 
derived  stannic  acid,  SnO(OII).2,  and  its  .salts,  as  SnO(ONa)2.  The  oxides  of  the  next 
fundamental  forms,  1^2^^  Ma^O^,  are  of  an  acid  nature.  The  corresponding  highest 
hydroxides  do  not  exist ; but  the  salts  of  the  anhydro-acids,  derived  from  them,  are  known  : 

IIBiOs  Ik^CrO^  HMnO^. 

Bismuthic  Chromic  Permanganic 

acid.  acid.  acid. 

In  comj)osition  these  acids  correspond  to  the  following  metalloid  acids: 

IINO3  IICIO,. 

Nitric  acid.  Sulphuric  acid.  Perchloric  acid. 

Salts  of  osmium  tetroxide  (OsO^)  are  not  known. 


OXIDES  AND  HYDROXIDES — HYDRATES. 


259 


T.ike  the  metalloids,  the  metals  of  the  last  four  series  form  lower  oxides  and  hydrates 
in  which  they  exhibit  a lower  valence  : 

SnO  BiaOs  MnO 

Sn(OH)2  Bi(OH)3  Mn(OH)2. 

These  lower  oxides  and  hydroxides  have  a basic  character,  and  in  their  whole  deport- 
ment resemble  the  corresponding  compounds  of  the  metals  of  the  first  three  groups. 


The  metals  of  the  first  two  groups  have  higher  oxygen  compounds, 
called  peroxides,  e.  g.,  Na202,  BaOj.  These  do  not  form  corresponding 
salts.  By  the  action  of  dilute  acids  hydrogen  peroxide  is  produced  : 

BaO,  -f  2HCI  = BaClj  -f  H^Oj. 

In  consequence  of  this  reaction,  it  is  very  probable  that  in  the  peroxides, 
the  oxygen  atoms  are  arranged  in  a chain-like  manner  as  in  hydrogen 
peroxide : 

Na  — O. 

Na  — O'^ 

Sodium  peroxide.  Barium  peroxide. 

When  concentrated  acid  acts  upon  them,  oxygen  is  evolved,  and  salts 
of  the  lower  oxides  result;  heated  with  hydrochloric  acid,  chlorine  is 
generated  : 

BaOj  + 4HCI  = BaClg  + 2H2O  T Clj. 

Ordinarily,  all  higher  oxides  which  evolve  chlorine  with  hydrochloric  acid  are  termed 
peroxides,  e.  g.,  Pb02,  lead  peroxide,  and  MnO.^,  manganese  peroxide.  However,  these 
latter  compounds  do  not  possess  the  structure  of  true  peroxides.  Lead  dioxide,  Bb02, 
is  wholly  analogous  to  tin  dioxide,  SnO.^,  and  is  capable  of  combining  with  bases ; it 
also  yields  a tetrachloride,  PbCb,  which  it  is  true  readily  breaks  down  into  the  dichloride 
and  chlorine  ; therefore  we  infer  that  the  two  oxygen  atoms  are  in  direct  union  with 
quadrivalent  lead.  So  manganese  is  probably  quadrivalent  in  manganese  peroxide. 
The  difference  between  these  oxygen  compounds  and  the  true  peroxides  is  shown  by 
their  inability  to  form  hydrogen  peroxide.  Persulphuric  and  percarbonic  acids  belong 
to  the  true  peroxides. 

Finally,  some  univalent  metals  are  capable  of  forming  oxides  contain- 
ing four  atoms  of  metal,  e.  g.,  K^O,  Ag^O  ; these  compounds  are  termed 
quadrant  oxides  or  suboxides. 


Salts. — By  the  action  of  bases  upon  acids,  salts  and  water  result : 
NaOH  + HNO3  = NaN03  + H2O, 

whereas  oxides  and  acid  anhydrides,  e.  g.,  CaO  and  CO2,  in  the  absolute 
absence  of  water  do  not  combine,  or  at  least  only  with  difficulty. 

d'hese  salts  are  also  i)roduced  like  the  halides  by  the  action  of  metals 
upon  the  acids : 


H2SO,  j Zn  = ZnSO,  f IL,  ; 


INORGANIC  CHEMISTRY. 


260 


most  of  the  oxygen  acids  suffer  simultaneous  reduction  liy  tlie  hydrogen 
evolved,  d'his  hapiiens  with  suli)huric acid  if  it  be  used  in  a concentrated 
form,  so  that,  e.  g.,  in  the  action  of  zinc  on  concentrated  .sul])huric  acid 
both  sulphurous  acid  and  hydrogen  sulphide  are  produced  ; 

II2  + 2II/); 

4ll2  + II,S  + 41120. 

Sulphurous  acid,  nitric  acid  and  chloric  acid  are  also  reduced  in  dilute 
solutions  (see  pp.  184,  186,  203).  Phosphoric  acid,  in  contrast  to  arsenic 
and  antimonic  acids,  is  very  stable  in  the  presence  of  nascent  hydrogen 
( pp.  148,  216).  The  salts  of  oxygen  acid,  like  those  of  the  haloid  acids,  are 
to  be  regarded  as  derived  from  the  acids  by  replacement  of  their  hydrogen 
by  metals.  The  acids  themselves  can  be  viewed  as  hydrogen  salts.  'This 
was  the  conception  of  IT  Davy  (1815),  of  P.  I..  Dulong  (1816),  and  par- 
ticularly of  J.  Liebig  (1838),  from  chemical  considerations,  and  also  a 
little  later  (1839)  of  J.  Fr.  Daniell,  from  electrolytic  reasons. 

As  we  have  seen,  the  poly  basic  acids  yield  the  primary,  secondary,  ter- 
tiary, etc.g  salts  by  the  replacement  of  one  or  several  hydrogen  atoms  by 
metal  (see  p.  172).  In  the  same  manner  n-series  of  salts  are  derived  from 
n acid  bases  (which  contain  n-OH  groups)  : 


.OH 

rOH 

/OH 

(NO, 

Bi^OH 

OH 

BifN03 

Bi  \ NO, 

^OH 

iN03 

\no: 

[no,. 

Such  salts  in  which  not  all  the  hydroxyl  groups  of  the  polyacid  hydrox- 
ide are  replaced  by  acid  residues  are  called  basic : 


Basic  lead  nitrate.  Basic  zinc  chloride. 


Besides  these  basic  salts  there  exist  some  of  another  form.  We  saw 


that  the  polybasic  acids  can  combine  to  form  poly-  or  anhydro-acids ; 
similarly,  the  polyhydric  bases  form  polyhydrates : 


SO,<OH 

S02<0H 


Cu<OH 

Cu<OH 


Pb< 

Pb< 

Pb< 


OH 

O 

O 

on, 


from  which  basic  salts  are  obtained  (see  Copper  and  Lead)  by  replace- 
ment of  hydroxides  by  acid  residues. 

By  the  replacement  of  the  hydrogen  atoms  in  the  polybasic  acids  or  of 
the  hydroxyl  groups  in  the  polyhydric  bases  by  various  radicals  we  get  the 
so-called  mixed  or  double  salts : 


K 

NIL 

H 


Potassium  ammonium 
phosphate. 


f‘<Cu 

SO.<K 


Potassium  copper 
sulphate. 


S0,< 

so,^ 


K 

A1 


Potassium  aluminium 
sulphate. 


’‘''<co 

I’l><c|  " 


MgClj.  KCl 


NO3 

Cl, 2 


AuCb.KCl  PtCh.2KCl. 


ACTION  OF  METALS  UPON  SALTS  AND  ACIDS. 


261 


Just  as  the  fluorides  of  boron  and  silicon,  BFI3.  KFl,  SiFl^.  2KFI,  are 
derived  from  peculiarly  constituted  atomic  acids,  HBFl^,  H.^SiFl^^  (pj). 
170,  226),  so  also  many  of  the  so-called  double  chlorides,*  e.  g.,  PtCl^.- 
2KCI,  must  be  viewed  as  atomic  in  their  structure  (see  p.  271  and  under 
Platinum). 


ACTION  OF  METALS  UPON  SALTS  AND  ACIDS. 

We  have  seen  that  the  metals  by  solution  in  acids  are  able  to  form  salts, 
h\  drogen  being  simultaneously  evolved. 

The  metals  deport  themselves  in  the  same  manner  with  the  salts.  Zinc 
introduced  into  a solution  of  copper  sulphate  is  dissolved  to  sulphate  and 
metallic  copper  is  deposited  : 

Zn  -f-  CuSO^  = ZnSO^  -f-  Cu. 

Herein  is  shown  the  perfect  analogy  between  acids  and  salts : 

Zn  + H^SO,  = ZnSO,  + H^. 

In  chemical  nature  hydrogen  is  a metal.  Hence  the  acids  maybe  viewed 
as  hydrogen  salts. 

The  displacement  of  a metal  from  its  salts  by  others  appears  to  be  in- 
fluenced by  its  electrical  deportment  in  so  far  as  it  occurs,  in  general,  in 
accordance  with  its  position  in  the  electric  tension  series.  Indeed  the 
more  electro-positive  (basic)  metals  replace  the  electro-negative  (less 
basic).  In  the  following  series  each  metal  throws  out  from  solution  those 
preceding  it : Au,  Pt,  Ag,  Hg,  Cu,  Pb,  Sn  (Fe,  Zn).  Iron  and  zinc  pre- 
cipitate almost  all  the  heavy  metals  from  solutions  of  their  salts.  The 
strongly  positive  potassium  and  sodium  are  able  to  displace  all  other 
metals.  This  is  very  evident  from  their  action  upon  fused  haloid  salts — 
a reaction  which  frequently  serves  for  the  separation  of  the  metals  in  a 
free  condition  : 

AICI3  + sNa  = A1  + sNaCl. 

In  its  electrical  deportment  hydrogen  stands  near  zinc  ; like  the  latter,  it 
must,  therefore,  displace  all  more  negative  metals.  If  this  does  not  happen, 
the  cause  must  be  sought  in  the  volatility  of  hydrogen ; in  fact,  we  know 
that  hydrogen,  under  powerful  pressure,  is  capable  of  separating  gold, 
silver  and  some  other  metals  from  their  salt  solutions. 

In  addition  to  electrical  deportment,  the  nature  of  the  solution  influences  the  precipi- 
tation of  one  metal  by  another.  Thus,  lead  separates  tin  from  its  chloride,  wSnCq,  while, 
on  the  other  hand,  tin  throws  out  lead  from  the  solution  of  its  oxide  in  alkalies.  This 
recalls  the  fact  that  iodine  is  displaced  in  hydriodic  acid  and  the  metallic  iodides  by  chlo- 
rine ; and  that  conversely  iodine  replaces  chlorine  in  its  oxygen  derivatives  (p.  278). 

Mylius  and  Fromm  (Her.  27  (1894)  i,  630)  claim  that  when  positive  metals  act  upon 
dilute  salt  solutions  of  negative  metals,  alloys  are  produced,  because  in  the  moment  of 
their  precipitation  metals  are  capable  of  uniting  at  the  ordinary  temperature.  Generally 
these  alloys  are  y)orous,  l)lack,  and  either  amorphous  or  crystalline.  The  following  crys- 
talline compounds  were  obtained  : Cu.^Cd,  AuOkj,  CiqSn  and  lead-platinum.  Zinc  pre- 


* Consult  American  Chemical  Journal,  li,  p.  291. 


262 


INORGANIC  CHKMIS'I’RV. 


cipitates  a zinc- silver  alloy  from  dilute  silver  solutions,  which  is  converted  into  white  crys- 
talline silver  by  a concentrated  silver  solution,  and  also  hy  strong  acids  : 

Ag„Zn  I 2AgN()3  = Ag(„  t 2)  + /^(NOa)^. 

The  other  alloys  are  similarly  decomposed  hy  acids,  the  negative  element  being  set  free. 


ELECTROLYSIS  OF  SALTS. 

Conductors  are  divided  into  two  classes,  depending  upon  their  behavior 
in  the  passage  of  the  electric  current.  Conductors  of  the  first  class  are 
not  decomposed  in  the  passage  of  the  current : the  metals,  peroxides  and 
carbon.  Conductors  of  the  second  class  carry  the  current  with  simul- 
taneous decomposition  : “ Electrolysis  ” — salts,  acids  and  bases,  fused  or 
in  aqueous  solution.  They  are  also  termed  electrolytes  (p.  92).  In 
most  instances  the  conductivity  of  the  first  class  diminishes  with  rise  in 
temperature,  while  it  increases  with  the  second  class. 

On  subjecting  a salt  in  a fused  or  dissolved  condition  to  the  action  of 
an  electric  current,  it  is  decomposed,  so  that  the  metal  separates  at  the 
negative  pole  and  the  acid  group  or  halogen  at  the  positive : 

NaCl  = Na  4 Cl 
CuSO,  = Oi  + SO,. 


The  liberated  acid  residues,  e.  g.,  SO,,  cannot  exist  in  a free  condition  ; 
they  break  down  into  oxygen  and  an  acid  oxide,  which,  with  the  water 
of  the  solution,  again  forms  the  acid  : 

SO,  -f  H^O  = H^SO,  + O. 

Thus,  in  the  electrolysis  of  such  salts,  the  metal  and  oxygen  separate 
out — the  former  at  the  negative,  the  latter  along  with  free  acid  at  the 
])ositive  pole  (see  p.  92  note). 

All  neutral  salts  are  similarly  decomposed.  If,  however,  the  metal 
contained  in  the  salt  acts  upon  water  when  free,  manifestly  a secon- 
dary reaction  must  occur  at  the  negative  pole.  The  real  electrolytic 
decomposition  of  potassium  sulphate  would  then  take  place  according  to 
the  following  equation  : 

K3SO,  = K,  + (S^,). 

d’he  sei)arated  potassium  decomposes  the  water  with  formation  of  potas- 
sium hydroxide  and  the  disengagement  of  hydrogen  : 

2K  -f  211,0  2K0II  + 11.,. 

Therefore,  hydrogen  and  potassium  hydroxide  occur  as  definite  decom- 
j)osition  ])roducts,  at  the  negative  ])ole  (the  kathode);  at  the  positive 
(the  anode),  however,  we  have  oxygen  and  sulphuric  acid.  On  coloring 
the  liquid  exposed  to  the  electrolysis  with  a little  violet-syrup,  that  part 


ELECTROLYSIS  OF  SALTS. 


263 


at  the  positive  pole  will  become  red,  owing  to  the  acid  formed,  while  that 
at  the  negative  pole  will  be  colored  green  by  the  base.  That  the  electro- 
lytic decomposition  of  potassium  sulphate  and  similar  salts  proceeds  in 
the  manner  given,  may  be  proved  experimentally  by  using  mercury  as  the 
negative  electrode ; then  the  separated  potassium  will  combine  with  the 
mercury  and  form  an  amalgam,  which  will  act  gradually  upon  the  water. 

The  bases  and  acids  are  decomposed  in  this  manner  by  the  current. 
Thus,  as  Davy  (1807)  showed,  molten  caustic  potash,  KOH,  breaks  down 
into  K and  OH;  the  first  separates  in  metallic  form  upon  the  negative 
pole  (and  gradually  acts  upon  the  fused  KOH  with  the  liberation  of 
hydrogen),  while  at  the  positive  pole  water  and  oxygen  appear — produced 
from  the  OH-ions : 

2(0H)  = H^O  + O. 

Acids  when  electrolyzed  behave  like  hydrogen  salts.  Hydrochloric 
acid,  for  example,  breaks  down  into  hydrogen  and  chlori-ne.  Sulphuric 
acid  (depending  upon  its  strength)  is  decomposed  in  aqueous  solution 
into  the  ions  2H  and  SO^  or  H and  HSO^.  The  anion  SO^  (or  HSOJ 
is  immediately  converted  by  the  water  into  sulphuric  acid  and  oxygen  : 

2HSO,  + 2H2O  = H2SO,  -h  O. 

However,  under  certain  conditions,  an  appreciable  amount  of  persul- 
phuric  acid  is  produced  : 

2HS0^  = H.^S^Og. 

[See  p.  188  and  Richarz,  Ber.  21  (1888),  1673.] 

The  electrolysis  of  acid  sulphates  and  of  acid  carbonates  can  also  pro- 
ceed in  the  direction  last  indicated;  they  break  down  into  kations  H 
and  anions  MeSO^  and  MeCOg,  from  which  arise  persulphates  and  per- 
carbonates:  MegSgOg  and  MegCgOg ; see  also  p.  231. 

In  the  thirties  of  the  present  century  Michael  Faraday  demonstrated  that 
when  a compound  is  decomposed  by  the  electric  current  the  quantity  of 
the  material  so  broken  down  is  proportional  to  the  quantity  of  electricity 
which  has  passed  through  it.  He  also  discovered  the  following  law  bear- 
ing his  name : 

When  the  electric  cu7'rent  passes  through  differefit  deco7nposable  conduc- 
tors a7'ranged  m series^  the  qua7itities  of  the  substances  separated  si77iulta- 
77eously  are  to  each  other  as  the  ratio  of  their  che7nical  equivale7its. 

Thus  in  the  simultaneous  decomposition  of  hydrochloric  acid,  dilute 
sulphuric  acid  and  ammonia*  (pp.  77,  98),  equal  volumes  of  hydrogen  are 
liberated,  while  at  the  positive  pole  i volume  of  chlorine,  volume  of 
oxygen,  and  ^volume  of  nitrogen  appear,  after  the  solutions  have  become 
saturated  with  the  gases.  The  quantities  decomposed  by  electrolysis, 
therefore,  bear  the  following  relation  : 

H0SO4  NH3 
’ 2 ’ 3 • 


*The  hydrochloric  acid  sliould  contain  219  grains  in  a liter,  d'his  can  be  olitained  by 
mixing  200  grams  of  fuming  hydrochloric  acid,  of  specific  gravity  1.185,  with  128  grams 
of  water.  P'or  the  electrolysis  of  ammonia  use  i volume  of  concentrated  ammonia  water 
with  12.5  volumes  of  a saturated  salt  solution. 


264 


INORGANIC  CHEMISTRY. 


In  tlie  same  way,  equal  quantities  of  chlorine  are  set  free  from  all  metallic 
chlorides,  while  the  (juanlities  of  the  j)recipitated  metals  agree  with  the 
values  according  to  which  they  enter  chemical  action,  their  equivalent 
weight.  The  quantities  of  the  different  salts,  decomposed  ])y  electrol- 
ysis, stand  in  the  following  relation  : 

^ Cu^l-J  ShCla  FcClo  FeCls  S11CI2  S11CI4  CuSOj  Hk(CN)2  Hfi:2(N03)2 

”'^*2'2*3*2*3*2*4*  2*  2 * 2 ' 

Therefore,  31.8  parts  of  cojiper  are  deposited  for  the  35.45  parts  of 
chlorine  in  cupric  chloride,  CuCl.^,  but  from  cuprous  chloride,  Cu./'l^, 
we  obtain  63.6  parts  of  copper;  from  mercuric  cyanide,  Hg(CN).^,  we 
obtain  100.15  parts  of  mercury,  and  from  mercurous  nitrate,  }ig(N03), 
200.3  parts  of  mercury,  etc. 

Tliese  relations  can  be  ol)litcratecl  in  that  reductions  of  the  electrolyte  at  the  negative 
pole  and  its  oxidation  at  the  positive  may  occur  through  the  decomposition  products. 
Thus,  in  the  electrolysis  of  cuprous  chloride,  cupric  chloride  instead  of  chlorine  will  be 
found  at  the  anode  : 

CuCl  4-  Cl  = CuCh, 

and  in  the  electrolytic  decomposition  of  cupric  chloride  cuprous  chloride  will  appear  at 
the  kathode  : 

CuCb  + Cu  = 2CuCl. 

Some  obscurity  of  the  ratios  acquired  by  the  law  also  arises  if  the  salt  is  hydrolytically 
decomposed,  e.g.,  if  water  converts  it  partly  into  hydroxide  or  basic  salt,  and  free  acid  : 

FeCh  + 3H,0  = Fe(OH)3  + 3110  ; 

the  free  acid  is  also  subject  to  electrolytic  decomposition.  Further,  the  proportions 
change  iCthe  electrodes  are  attacked  by  the  substances  liberated  upon  them  (p,  77). 
For  example,  if  in  the  electrolysis  of  copper  sulphate  the  positive  electrode  consists  of 
copper  then  oxygen  will  not  be  liberated  here  (SO^  = SO3  -}-  O),  but  just  as  much 
metallic  copper  will  dissolve  as  separates  upon  the  kathode  : 

Cu  + SO4  = CuSO, ; 

consequently  the  copper  content  of  the  solution  remains  unchanged. 

H.  Helmholtz  accordingly  expresses  Faraday’s  law  as  follows; 

The  same  quantity  of  electricity  acting  upon  different  electrolytes  liberates 
equal  valences  or  combines  the7n  in  sojue  new  way. 

Causs  and  Weber  referred  most  electrical  and  magnetic  magnitudes  to  length  (cm.), 
mass  (gram)  and  time  (sec.). 

'I'he  units  derived  in  this  way  are  the  absolute  measurements.  The  unit  quantity  of 
electricity  in  this  measurement  precipitates,  aecording  to  Kohlrausch,  0.01118  gram  of 
silver.  'I’he  tenth  part  of  this  quantity,  or  that  which  will  precij)itate  0.001 1 18  gram  — 1. 1 18 
mg.  of  silver,  has  been  selected  as  the  “practical  unit”  and  is  called  the  “coulomb” 

(coul.)  ; 300  mg.  of  silver  deposited  iq^on  the  kathode  indicate  therefore  that  ^^^^^  = 268 
coni,  have  passed  through  the  solution.  If  this  took  place  in  ten  minutes,  then  in  one 
second  ^ = 0.447  coul.  would  have  passed  through  the  liejuid.  'I'he  current  strength 
would  then  have  been  0.447  miipere.  'I'he  current  strength  (amp.),  therefore,  represents 


SOLUTIONS. 


265 


the  number  of  coulombs  which  pass  through  the  entire  cross-section  of  the  current’s 
course.  A current  of  one  ampere  deposits  or  separates  : 


Mg.  Silver.  Mg.  Copper. 
In  one  second,  . . 1.118  0.3284 

In  one  minute,  . . 67.08  i9-7o 

In  one  hour,  . . 4025  1182 


C.c.  Electro- 
lytic Gas  C.c.  Hydrogen 


AT  0°  AND  AT  0°  AND 

Mg.  Water.  760  Mm.  760  Mm. 

0.0933  0.1740  o.  n6o 

5.60  10.44  6.960 

335-9  626  417 


The  electro-chemical  equivalents  are  the  quantities  whicli  are  separated  out  or  converted 
into  other  compounds  in  the  unit  of  time.  It  is  possible  in  accordance  with  Faraday’s 
law  to  calculate  readily  from  the  silver  and  copper  values  those  of  other  elements 
or  atomic  groups.  In  general,  a current  of  i ampere  deposits  in  i sec.  0.01036,  in  i 

min.  0.6215,  ^ hour  37.29  mg.  — equivalents  of  any  substance  ~ 

0.01036  ; Ag  = 107.93  ; “ — 31-^)  » hence,  35.45 . 0.01036  — 0.3673  mg.  chlorine  in  a 
second,  etc.  It  may  also  be  mentioned  that  the  practical  unit  of  resistance,  the  ohm,  is  a 
column  of  mercury  106.3  cm.  in  length  and  i .square  mm.  in  section,  and  that  the  volt,  the 
practical  unit  of  electromotive  force,  is  the  power  which  in  a closed  circuit  of  an  ohm 
resistance  produces  a current  of  i ampere. 


Faraday’s  law  requires  that  when  equal  quantities  of  electricity  pass  through  different 
electrolytes,  the  same  number  of  valences  must  be  liberated  or  converted  into  new  com- 
pounds at  each  pole.  Therefore,  the  decomposition  of  aqueous  hydrochloric  acid  is 
accomplished  by  the  consumption  of  the  same  quantity  of  electricity  which-  would  be  re- 
quired by  equivalent  quantities  of  hydrobromic  acid,  hydriodic  acid,  etc.  We  know  on 
the  other  hand  that  very  different  amounts  of  heat  are  necessary  to  effect  the  decompo.si- 
tion  of  equivalent  quantities  of  these  compounds.  Hence  we  might  expect  that  their 
electrolytic  decomposition  would  demand  a varying  expenditure  of  electricity.  The 
endothermic  bodies  should  be  most  easily  decomposed,  while  those  with  the  greater  heat 
of  formation  and  heat  of  decomposition  would  be  most  difficult. 

Furthermore,  so  long  as  the  electric  force,  active  in  the  conductor,  is  not  equal  in 
strength  to  the  affinity,  decomposition  does  not  occur,  but  when  these  forces  are  equal, 
then  many  molecules  will  at  once  be  broken  down.  This  contradicts  all  experience 
which,  according  to  R.  Clausius,  may  be  explained  as  follows  [Fogg.  Annalen  (1857) 
loi,  338;  die  mechanische  Warmetheorie  (1879)  Ed.  ii,  164]  : The  current  does  not 
decompose  the  electrolyte  into  its  constituents.  In  consequence  of  the  heat  motions  of 
its  molecule  the  electrolyte  is  already  partially  dis.sociated  and  the  current  merely  separates 
the  already  free  ions  spacially.  Williamson,  at  an  early  date,  had  expre.s.sed  a similar 
view  in  an  unusually  able  paper  on  the  theory  of  the  ether-formation  [Ann.  Chem. 
Fharm.  (1851)  77,  37].  He  thought  that  in  a mass  of  molecules  the  latter  continually 
exchanged  the  atoms  and  atom  groups,  of  which  they  consisted,  with  one  another.  For 
example,  in  hydrochloric  acid  the  molecule  of  hydrogen  chloride  did  not  constantly  con- 
sist of  the  same  hydrogen  atom  and  chlorine  atom,  but  that  there  was  rather  a mutual 
exchange  of  corre.sponding  atoms  among  the  molecules.  S.  Arrhenius  has  in  late  years 
expanded  and  developed  these  ideas  into  the  theory  of  electrolytic  dissociation,  which 
will  receive  attention  later,  p.  268. 


SOLUTIONS. 

Most  liquids  can  take  up  gases,  other  liquids  and  solids  in  such  a man- 
ner that  a new  liquid,  homogeneous  both  physically  and  chemically — a 
solution — is  produced  (p.  256).  No  deep-seated  chemical  changes  occur 
between  the  solvent  and  the  substance  dissolved,  otherwise  the  process 

23 


266 


INORGANIC  CHEMISTRY. 


could  not  be  defined  as  solution  in  the  sense  indicated.  If,  for  example, 
carbon  dioxide  is  conducted  into  and  absorbed  by  caustic  potash,  a solu- 
tion of  carbon  dioxide  does  not  result,  but  rather  one  of  potassium  car- 
bonate. In  the  same  way  we  cannot  say  that  a metal  dissolves  in  an 
acid,  because  it  is  first  changed,  with  the  evolution  of  hydrogen,  into  a 
salt  which  dissolves.  Gases  can  mix  to  any  degree  with  one  another,  but 
the  solvent  power  of  liquids  for  gases  or  other  liquids  or  solids  is  gener- 
ally limited  and  also  dependent  upon  jiressure  and  temperature.  Aqueous 
solutions  are  alone  important  here,  and  only  they  will  be  considered 
now. 

The  absorption  of  gases  (Henry,  1803)  or  gas  mixtures  (Dalton,  1807), 
as  mentioned  on  p.  229,  is  proportional  to  tlie  pressure  under  which  the 
gas  exists.  The  gas  dissolved  by  a liquid  can  be  expelled  from  it  by 
raising  the  temperature,  by  lowering  the  pressure,  or  by  introducing 
another  gas,  e.  g.,  carbon  dioxide,  by  passing  air  through  the  solution,  or 
by  shaking  the  latter  with  air.  Certain  gas  solutions  are  distinguished 
by  the  fact  that  under  a definite  pressure  they  will  distil  without  change 
in  their  composition  at  a constant  temiierature ; (.see  Hydrochloric  Acid, 
p.  59,  Hydrobromic  Acid,  p.  62). 

Many  metallic  salts  separate  in  union  with  water — water  of  crystalliza- 
tion— from  their  solutions.  This  can  frequently  be  noticed  in  their 
change  of  color.  Anhydrous  copper  sulphate,  CuSO^,  is  white,  while 
its  aqueous  solution  is  blue,  and  on  evaporation  blue  hydrous  CuSO^ . sH^O 
crystallizes  out.  This  salt  loses  a portion  of  its  water,  on  exposure  to 
the  air,  at  the  ordinary  temperature — it  effloresces.  It  becomes  anhydrous 
above  200°.  Hence,  we  can  conclude  that  such  a hydrate  existed  in  the 
solution.  In  some  cases  the  metallic  salt,  depending  on  the  physical 
conditions  at  which  the  solution  or  crystallization  occurred,  unites  with 
varying  amounts  of  water,  which  is  recognizable  by  its  color.  The  com- 
pounds of  cobaltous  chloride,  CoCl.^,  with  one  and  two  molecules  of  water 
of  crystallization  are  blue  in  color  like  the  solution  from  which  they 
separate,  whereas  the  red  compound  with  six  molecules  of  water, 
CoClg  + 6H2O,  is  obtained  from  red-colored  solutions.  As  a rule, 
the  solubility  of  salts  increases  with  the  rise  in  temperature  and  is 
definite  for  every  temperature  {saturated  soliitiofi).  Thus,  i part  of 
potassium  nitrate  dissolves  in  4 parts  of  water  at  15°,  and  in  0.3-0. 4 
parts  at  the  boiling  temperature.  When  such  a hot  saturated  solution 
cools,  the  salt  crystallizes  out.  Sodium  chloride  is  a remarkable  excep- 
tion to  this,  as  it  dissolves  very  little  more  in  hot  than  in  cold  water: 
100  ])arts  of  water  at  the  ordinary  temperature  take  up  36  parts,  and  when 
boiling  39  jiarts  of  the  salt.  The  solubility  with  other  salts  increases  up 
to  a definite  temperature  and  then  it  diminishes.  It  is  thus  with  sodium 
sulphate,  Na.^SO^ ; 100  parts  of  water  at  0°  dissolve  five  parts  of  the  salt, 
and  the  most,  namely  55  jiarts,  at  34°  ; at  60°  only  45  and  at  100°  about 
42  parts  (see  Glauber’s  Salt).  When  a salt  dissolves  it  absorbs  heat  and 
the  water  cools.  However,  if  the  salt,  like  anhydrous  copper  sulphate 
or  calcium  chloride,  (kaCl.^,  can  take  u])  water,  i.  e.,  able  to  form  a 
hydrate,  then  heat  is  evolved  in  its  solution,  because  the  quantity  of 
heat,  set  free  in  the  formation  of  the  hydrate,  exceeds  that  necessary  to 


SOLUTIONS.  267 

dissolve  it.  Therefore,  hydrated  calcium  chloride,  CaCl^  -f  6H2O,  dis- 
solves in  water  with  the  lowering  of  temperature. 


THEORY  OF  DILUTE  SOLUTIONS. 

Many  chemists  regard  solution  as  a chemical  process  induced  by  the  chemical  attrac- 
tion between  the  solvent  and  the  substance  dissolved,  and  leading  finally  to  very  unstable 
and  undetermined  chemical  compounds  (Berthollet,  Mendelejeff,  Berthelot),  while  others 
believe  it  to  be  physical,  e.  g.,  Dossios,  who  considers  it  a sort  of  diffusion,  which  may  be 
compared  in  every  respect  to  vaporization.  Neither  of  the  two  views  alone  is  sufficient 
to  explain  satisfactorily  the  process  of  solution  [see  Z.  f.  anorg.  Ch.  6 (1894)  392]. 

While  we  cannot  be  certain  as  to  the  forces  which  bring  about  solution,  yet  since 
1885  J.  H.  van’t  Hoff  has  by  his  pioneer  researches  shown  that  the  condition  of  a sub- 
stance in  dilute  solution  is  similar  to  the  gas  condition.  The  equation  deduced  on 
page  12 1 : 

pv  = RT 


embodies  the  law  of  gases  (Boyle  and  Gay-Lussac)  and  that  of  Avogadro,  and  as 
proved  by  van’t  Hoff  it  also  answers  for  dilute  solutions  if  “osmotic  pressure”  be 
substituted  for  gas  pressure  [Z.  f.  phys.  Ch.  i (1887)  487;  3 (1889)  198;  Ber.  27 
(1894)  6], 

Osmotic  pressure  is  the  pressure  exerted  against  a diaphragm,  separating  a solu- 
tion from  a solvent,  and  w'hich  is  permeable  to  the  solvent  but  not  to  the  dissolved 
substance  (see  p.  237).  Changes  brought  about  by  osmotic  pressure  play  a very 
important  role  in  the  living  organism  and  particularly  in  plants,  therefore  they 
were  first  more  accurately  studied  by  botanists — W.  Pfeffer  in  1877  and  H.  de  Vries 
in  1884. 

The  action  of  osmotic  pressure  may  be  demonstrated  as  follows  : A cylinder  of  about 
TOO  c.c.  capacity  is  filled  with  a syrupy  sugar  solution  and  then  made  air-tight  with  a 
covering  of  animal  membrane.  When  this  vessel  is  immersed  in  a vessel  of  water,  the 
membrane  gradually  forms  a curved  surface,  because  the  water  penetrates  through  it  to 
the  sugar  solution.  The  experiment  can  also  be  conducted  as  follows:  An  open 
cylinder  is  closed  at  one  end  with  a membrane  and  then  filled  with  a concentrated  sugar 
solution,  the  other  end  of  the  cylinder  being  provided  with  a stopper  which  carries  a 
narrow  glass  tube.  When  the  cylinder  is  immersed  to  the  stopper  in  water,  the  latter 
penetrates  the  membrane  and  as  a consequence  the  liquid  rises  gradually  in  the  tube  to 
a definite  height ; thus  a means,  though  not  very  exact,  is  afforded  for  the  measurement 
of  the  magnitude  of  the  osmotic  pressure. 

Pfeffer  and  de  Vries  made  measurements  of  osmotic  pressure  and  found  the  following 
relations  : 

1.  Osmotic  pressure  is  proportional  to  the  concentration  or  inversely  proportional  to 
the  volume  in  which  a definite  amount  of  substance  is  dissolved. 

2.  Osmotic  pressure  increases  with  constant  volume  in  proportion  to  the  absolute  tem- 
perature. 

3.  Quantities  of  dissolved  substances,  which  are  to  one  another  as  their  molecular 
weights,  exert  equal  osmotic  pressure  when  dissolved  in  equal  volume  at  like  temperatures. 

One  of  van’t  Hoff’s  services  was  the  recognition  that  these  three  laws  of  osmotic  pres- 
sure correspond  to  the  laws  controlling  gases.  The  first  is  the  counterpart  of  Boyle’s 
law,  the  second,  of  the  law  of  Gay-Lussac  and  Dalton,  and  the  third  corresponds  to 
Avogadro’s  law. 

Let  us  apply  these  relations  to  a solution  which  contains  one  part  by  weight  of  cane 
sugar  in  lOO  parts  by  weight  of  water.  In  accordance  with  the  specific  gravity  of  such 
a solution  it  would  contain  i gram  of  sugar  in  100.6  c.c.  Pfeffer’ s experiments  show 
that  this  solution  exerts  at  0°  an  osmotic  pressure  which  would  hold  a mercury  column 
of  49.3  cm.  in  height  in  equilibrium.  If  it  is  desired  to  ascertain  whether  these  relations 
can  be  expressed  by  the  equation  p . v = 84800  T,  the  units  mentioned  on  p.  1 22  must  be 
chosen  and  the  pre.ssure  in  grams  per  square  centimeter,  the  volume  in  cubic  centimeters 


268 


inorc;anic  chemistry. 


for  the  gram-molecule  and  the  temperature  in  absolute  numbers  must  be  considered  in  the 
calculation.* 

In  this  particular  instance  the  value  for  p would  be  49.3  . 13.6  670.48  ; for  '1'  the 

value  273.  Cane  sugar,  CjjIIj./),,,  has  the  molecular  weight  342.18  ; tlie  gram-molecule 
of  the  solution  in  <|uesti()n  is  then  contained  in  342.18  . 100.6  — 34400  c.c.  (v)  approxi- 
mately. Hence  670.48  . 34400  - - R . 273  and  R 84500,  a value  which  accords 
within  the  limits  of  error  with  that  of  the  gas  constants. 

Conseciuently  dissolved  substances  exert  the  same  pressure,  as  o.smotic  pressure,  which 
they  would  exert  at  like  temperature  and  like  volume  in  the  gas  form.  'I’heir  mole- 
cules are,  as  in  the  case  of  gases,  .so  far  removed  from  one  another  that  only  the  proper- 
ties dej)en(lcnt  upon  their  number  manifest  themselves,  while  those  dej)ending  uix)n  com- 
position and  chemical  structure  ofttimes  disappear  absolutely.  Properties  of  this  kind, 
which  can  assume  like  values  for  chemically  com[)arable  quantities  of  the  most  different 
substances — h)r  molecules,  for  atoms  of  bodies  which  are  exceptions  to  Avogadro’s  law 
(P-  79)>  ions,  as  we  shall  .soon  .see,  are  termed,  following  Ostwald,  colligative proper- 
ties [co/ligare,  to  link  together).  The  f|uantities  of  two  substances,  which  in  dilute  .solu- 
tion or  in  the  gas  form  manifest  colligative  properties  of  ecjual  value,  will  be  to  one  an- 
other as  the  ratio  of  their  molecular  weights  if  the  exce])tions  indicated  be  not  considered. 
It  would  then  follow  : that  if  at  one  time  n-gram-molecules  of  a substance  A be  di.ssolved 
in  a definite  volume  of  a solvent,  and  then  n-gram-molecules  of  the  substance  P>,  in  both 
cases  certain  properties  of  the  original  .solvent  would  be  similarly  altered.  The  freezing 
point  will  fall  and  the  boiling  point  will  rise  regularly  in  both  instances  ; the  tension  of 
the  two  solutions  will  be  the  same  ; they  pos.sess  eipial  osmotic  jiressure,  ?.  e.,  they  are 
isotonic  {^Igoc,  equal  ; Tovn^,  tension).  It  is  in  this  manner  possible,  upon  comparing  solu- 
tions of  a substance  of  unknown  molecular  weight  with  those  of  one  of  known  molecular 
weight,  to  determine  the  unknown  molecular  weight:  (l)  by  isotony  (Pfeffer,  de  Vries)  ; 
(2)  by  lowering  of  tension  (Raoult)  ; (3)  by  the  rise  in  boiling  point  ( Reckmann-Arrhe- 
nius)  ; (4)  by  lowering  of  the  freezing  point  (van’t  IIoff-Raoult-Eykmann).  The.se 
methods  are  described  in  Richter’s  Organic  Chemistry.  To  show  how  their  properties 
change  in  proportion  to  concentration,  mention  may  be  made  of  the  following  : A .solution 
of  I part  of  common  salt  in  100  parts  of  water  begins  to  freeze  at  — 0.6°,  one  of  2 : 100 
at  — 1.2°,  of  4 : 100  at  — 2.4°,  of  14  ; too  at  — 8.4°,  i.  e.,  at  — 0.6  X Relow  the 

last-named  temperature  the  proportionality  does  not  hold  for  anhydrous  sodium  chloride, 
but  for  the  hydrate  NaCl  2H.pj.  In  the  case  of  other  .salts  the  hydrates  play  this  role 
at  higher  temperatures:  thus,  Nal  4H2O,  MgSO^  7H2O,  etc.  [Rudoidf  and  de 
Coppet]. 

THEORY  OF  ELECTROLYTIC  DISSOCIATION. 

The  laws  just  mentioned  serve  only  for  very  dilute  solutions,  just  as  the  laws  for  gases 
correspond  with  greater  accuracy  in  the  behavior  of  those  gases  which  are  far  removed 
from  the  vapor  condition  (p.  122).  And  even  very  dilute  solutions  show  deviations  from 
van’t  Hoff’s  theory  when  the  solvent  is  water  and  the  dissolved  .substance  is  an  electrolyte 
— a conductor  of  the  second  class  (p.  262).  The  salts,  acids  and  bases  are  electrolytes 
in  aqueous  .solution,  while  the  non-electrolytes  are  the  organic  compounds,  with  the  ex- 
cei)tion  of  the  pronounced  acids,  bases  and  salts  ; also  the  solutions  of  all  substances  in 
benzene,  carbon  bisulphide,  ether,  chloroform,  etc.  The  alcoholic  solutions  constitute  a 
transition  to  the  electrolytes. 

An  example  will  illustrate  the.se  important  relations.  While  the  aqueous  solutions  of 
ether,  glycol,  sugar,  urea  and  similar  non-conducting  organic  compounds,  which  contain 
the  molecular  weight  of  the  respective  sub.stances  expressed  in  grams  per  liter,  freeze  at 
— 1.8°,  the  freezing  point  of  the  corresponding  solutions  of  potassium  iodide,  sodium 
chloride  and  silver  nitrate  falls  to — 3-6°,  double  the  first  ; and  that  of  .such  a solution  of 
sodium  sulphate  to — 5.4°,  treble  the  first.  The  values  of  other  colligative  properties  of 
these  solutions  manifest  similar  variations  from  the  van’t  Hoff  theory  : rise  in  the  boiling 


* If  the  |)rcssiire  be  given  in  atmospheres  and  the  volume  in  liters,  then  R naturally 
acquires  another  value.  As  the  molecular  volume  (p.  98),  at  l atmosphere  and  0°, 
equals  22.4  liters,  then  R would  e(|ual  0.082  (because  p ::=  l,  v = 22.4,  T = 273)  ; 
lienee  in  general  [> . v 0.082  'f  (liter-atmosphere). 


SOLUTIONS. 


269 


point,  decrease  in  the  vapor  tension,  and  the  osmotic  pressure  is  also  double  or  treble 
that  required  by  the  theory.  Dilute  solutions  of  electrolytes  consequently  behave  as  if 
they  contained  more  molecules  of  the  dissolved  substance  than  the  corresponding  solu- 
tions of  non-electrolytes. 

Svante  Arrhenius’  theory  of  electrolytic  dissociation  [Z.  f.  phys.  Ch.  i (1887)  631  ; 
see  also  Planck,  ibid.  576]  offers  an  interpretation  for  this  abnormal  deportment. 
According  to  it  the  electrolytes — which  Hittorf  designates  salt-like  bodies  inasmuch  as 
the  acids  are  salts  of  hydrogen  and  the  bases  of  hydroxyl — do  not  exist  as  such  in  aqueous 
solution  but  have  broken  down  entirely  or  in  part  into  their  ions  (p.  265)  — NaCl  into 
+ — + — + — + — 

Na  and  Cl,  AgNOg  into  Ag  and  NO.^,  Na2S04  into  2Na  and  SO4,  and  KOH  into  K and  OH. 
The  anions  are  charged  with  negative  and  the  kations  with  positive  electricities  which,  in 
accord  with  the  law  of  Faraday,  are  equally  large  for  equivalent  quantities  of  different 
ions,  e.  g.,  the  ion  SO4  contains  twice  the  quantity  of  electricity  of  the  ion  Cl. 

The  ions  regarded  as  independent  mass-particles  for  comparison’s  sake  play  the  part  of 
molecules  and  accordingly  influence  the  values  of  the  colligative  properties.  This  would 
explain  the  exceptions,  cited  above,  to  the  theory  of  solutions;  the  molecule  NaCl, 
+ — _ -f 

resolved  into  the  ions  Na  and  Cl,  acts  like  two  molecules;  Na^vSO^  separated  into  2Na 

and  SO4  behaves  like  three  molecules  of  a non-electrolyte.  The  determination  of  the 
lowering  of  the  freezing  point,  etc.,  can  therefore  serve  in  determining  the  degree  of  dis- 
sociation. Dissociation  and  electric  conductivity  run  parallel  because  the  conduction  of 
the  current  is  influenced  only  by  the  presence  of  free  ions  and  their  quantity,  whereas 
molecules  which  are  not  dissociated  do  not  participate  in  conducting  the  current. 
Accordingly  the  conductivity  actually  increases,  if  molecular  quantities  be  considered, 
with  growing  dilution  and  attains  its  maximum  when  all  the  molecules  of  the  electrolyte 
are  dissociated. 

The  neutral  salts  are  most  strongly  dissociated  and  particularly  those  with  univalent 
ions,  e.  g.^  NaCl,  AgNOg,  KI,  NH4Br.  Usually  more  than  half  of  the  salt  is  present 
in  the  form  of  free  ions  in  aqueous  solutions  at  medium  concentrations.  Salts  with  poly- 
valent ions  are  less  completely  dissociated,  and  in  the  case  of  the  mercury  haloids  the 
dissociation  is  extremely  slight.  The  degree  of  dissociation  with  acids  and  bases  corre- 
sponds to  what  is  commonly  called  their  ‘‘strength”  ; the  strongest  bases  and  acids  are 
most  completely  dissociated.  A gradual  dissociation  occurs  with  the  polybasic  acids  : 

+ — 

thus  sulphuric  first  breaks  down  into  the  ions  H and  HSO4  ; the  univalent  anion  HSO4 

— -I- 

in  turn  dissociates  into  SO4  and  H. 

The  strong  acids  would  then  be : Hydrochloric,  hydrobromic,  hydriodic,  nitric, 
chloric,  perchloric  and  sulphuric,  as  well  as  the  polythionic  acids.  Strong  bases  : Alka- 
lies, alkaline  earths  and  the  oxide  of  thallium.  In  solutions  of  medium  concentration 
these  compounds  are  dissociated  more  than  half. 

Moderately  strong  acids:  Phosphoric,  suljDhurous  and  acetic.  Moderately  strong 
bases  : Ammonia,  magnesia  and  silver  oxide  (dissociation  does  not  exceed  10  per  cent.). 

Weak  acids  and  bases,  the  dissociation  of  which  is  in  part  scarcely  measurable  : Car- 
bonic acid,  hydrogen  sulphide,  hydrogen  cyanide,  silicic  acid,  boric  acid — the  hydroxides 
of  the  other  bivalent  and  trivalent  metals. 

The  formation  of  a salt  from  a base  and  an  acid  proceeds,  according  to  this  theory,  in 
aqueous  solution  as  follows  : The  base  breaks  down  into  the  metallic  kation  and  the 
anion  OH,  the  acid  into  th«  kation  H and  the  anion — the  acid  residue  j)reviously  com- 
bined with  it.  Because  water  j)ossesses  an  extremely  slight  degree  of  dissociation,  being 
a non-conductor  (a  gram  equivalent  of  its  ions  is  contained  in  about  10,000,000  liters), 
it  will  be  produced  whenever  the  ions  H and  OH  meet,  so  that  in  the  present  case,  if  we 
have  started  with  equivalent  quantities,  only  the  ions  of  the  salt  will  remain  in  solution  : 

Na  + OH  -f  Cl  + H = Na  + Cl  + H^O + 13-7  Cal. 


K + OH  + N^  +^H  = K -f  NO3  -f  H2O -f  13.7  Cal. 

Ca  4-  2OH  j 2CI  -f  211  = Ca  -f  2CI  4-  2H2O 4-  27.4  Cal. 


270 


INORGANIC  CHEMISTRY. 


'I'he  common  and  essential  thing  in  these  changes  is  the  prodnctioti  of  water  from  the 
ions  11  and  (Jll  ; the  other  ions  remain  unchanged  as  long  as  the  concentration  of  tile 
solution  is  not  altered.  A fact  in  liarmony  with  this  is  that  upon  fieufraiizin^  equivalent 
quantities  of  strong  bases  with  strong  acids  an  equal  heat  modulus  (thermal  value)  is 
obtained ; hence, 

(II  acp,  OH  aq.)  = 13.7. 

This  fact,  for  which  no  reason  could  be  previously  observed,  now  seems  in  a similar  man- 
ner to  be  demanded  by  and  also  to  conhrm  the  theory  just  discussed. 

d'he  salt-like  bodies  according  to  the  theory  of  electrolytic  dissociation  appear  to  be 
binary  in  their  constitution.  Berzelius’  old  electro  chemical  theory  also  made  them  con- 
sist of  electro-positive  and  electro  negative  parts — they  were  dualistically  constituted. 
The  salts  of  the  oxygen  acids  consisted  of  the  acid  anhydride  and  the  metallic  oxide,  e.g., 
— 

potassium  sulphate  of  SO.,.  — whereby  they  were  brought  in  opposition  to  the  haloid 

salts,  which  facts  did  not  justify,  d'he  electro-chemical  theory  of  Arrhenius  proceeding 
from  positive  and  negative  constituents  has  a dualistic  character,  attaches  itself  thereby  to 
the  idea  of  Daniell  and  of  Liebig  (p.  260),  and  corresponds  to  the  views  to  which  Wil- 
liamson and  Clausius  were  led,  the  former  for  chemical  and  the  latter  for  electrolytic 
reasons  (p.  265).  d'his  theory  at  first  appeared  strange  and  met  great  oj^position,  but 
it  is  now  almost  universally  accepted  and  it  is  generally  admitted  that  we  are  indebted  to 
it  for  many  great  and  surprising  discoveries  in  the  domain  of  chemistry  and  physics  which 
had  gone  unobserved.  It  appears  almost  absolutely  necessary  for  young  chemists  to 
make  themselves  conversant  with  this  theory  and  to  that  end  the  publications  cited  on 
pp.  49,  66  will  be  found  helpful.  The  following  will  also  prove  especially  valuable  : 
( Istwald,  Foundations  of  Analytical  Chemistry  ; Liipke,  Elements  of  Electro  chemistry  ; 
Le  Blanc,  Electro-chemistry ; Lob,  The  Elements  of  Electro  chemistry  ; Windisch, 
Determination  of  the  Molecular  Weights. 


TRANSPOSITION  OF  SALTS. 

When  two  salts  in  solution  or  fusion  come  together,  a chemical  action  will  frequently 
occur.  Claude  Louis  Berthollet  (Es.sai  de  statique  chimique,  1803)  endeavored  to  ex- 
plain the  resulting  phenomena  by  referring  them  to  purely  physical  causes,  and  excluded 
chemical  affinity. 

In  the  opinion  of  Berthollet,  four  salts  always  arise  in  the  solution  of  two.  For 
example,  on  mixing  solutions  of  copper  sulphate  and  sodium  chloride,  there  exist  in  solu- 
tion copper  sulphate,  sodium  sulphate,  copper  chloride,  and  sodium  chloride  : 

nCuSO^  -j-  mNaCl  yield 

(n  — x)CuSO„  (m  — 2x)NaCl  -)-  xNa2SO„  -j-  xCuCl2. 

That  copper  chloride  is  really  present  in  the  solution  together  with  the  sulphate, 
follows,  from  the  fact  that  the  blue  color  of  the  latter  acquires  a greenish  color,  peculiar 
to  the  copper  chloride,  by  the  addition  of  sodium  chloride  ; other  phenomena  are  not 
noticeable  at  first.  Suppose  one  of  the  four  salts  formed  in  the  solution  is  insoluble  or 
volatile,  the  reaction  will  occur  somewhat  differently.  Upon  adding  barium  chloride  to 
the  copper  sulphate  solution  four  salts  will  be  formed  at  the  beginning  just  as  in  the  first 
case.  The  barium  sulphate  produced  separates,  however,  in  consequence  of  its  insolu- 
bility, the  ecjuilibrium  of  the  four  salts  will  be  disturbed,  and  new  quantities  of  copper 
sulphate  and  barium  chloride  act  upon  each  other  until  the  transposition  is  complete  : 

CuSO,  -f  BaCl2  = BaSO^  -f  CuCl.,. 

'I'he  chemical  transposition  may,  therefore,  be  explained  by  the  insolubility  of  the  barium 
sulphate.  On  adding  hydrochloric  acid,  or  soluble  chlorides,  to  the  solution  of  a silver 
salt  all  the  silver  is  precipitated  as  chloride,  becau.se  the  latter  is  in.soluble. 

'I'ake  another  example.  On  adding  sulphuric  acid  to  a solution  of  potassium  nitrate 
there  is  apparently  no  pcrce])tible  alteration  ; yet  four  compounds,  KNO3,  KIISO4,  H2SO4 
and  UNO,,,  are  [)re.sent  in  tlie  .solution.  This  was  proved  by  the  thermo-chemical  invest!- 


TRANSPOSITION  OF  SALTS. 


271 


gationsof  Julius  Thomsen,  and  from  determinations  made  by  W.  Ostwald  on  the  changes 
in  volume  and  density  which  are  connected  with  transpositions.  The  two  acids  distribute 
themselves  upon  the  base.  The  proportion  or  degree  to  which  this  occurs  is  dependent 
upon  the  quantity  of  potassium  nitrate  and  sulphuric  acid  in  a volume  unit,  from  external 
circumstances,  such  as  the  temperature,  and  upon  the  nature  of  reacting  substances.  The 
more  sulphuric  acid  there  is  in  proportion  to  the  nitric,  the  more  sulphate  will  there  be 
formed.  This  is  a case  of  mass-action^  the  theory  of  which  was  developed  by  Guldberg 
and  Waage  (p.  95).  On  heating  or  evaporating  the  nitrate  solution  containing  sulphuric 
acid  a new  condition  arises  : the  volatility  of  nitric  acid.  Hence  it  follows  that  on 
evaporation  the  transposition  proceeds  to  completion  in  the  sense  of  the  equation  : 

2KNO3  -f  2H2SO4  = KHSO^  + 2HNO3. 

These  relations  are  thus  explained  by  the  theory  of  electrolytic  disso- 
ciation. Neutral  salts  in  dilute  solutions  do  not  as  a rule  act  upon  one 
another  because,  like  the  salts  which  can  arise  from  them  by  transposition, 
they  are  equally  dissociated.  Thus  a dilute  solution  of  potassium 
chloride  contains  the  ions  K and  Cl,  that  of  sodium  nitrate,  the  ions  of 
Na  and  NO,.  If  these  solutions  are  mixed  no  change  occurs.  A solu- 
tion of  exactly  the  same  properties  would  be  obtained  on  mixing  corre- 
sponding amounts  of  potassium  nitrate  and  sodium  chloride. 

If  a substance,  less  soluble  and  less  dissociated  under  the  prevailing  conditions,  can 
be  formed  from  the  ions,  then  a transposition,  apparently  an  exchange,  will  occur.  It  is 
in  this  fashion  that  the  transpositions  described  above  are  to  be  explained.  The  neutral- 
ization of  bases  by  acids  is  to  be  accounted  for  in  this  way  : water,  which  is  but  slightly 
dissociated,  is  formed  from  the  ions  H and  OH,  whereas  the  ions  which  belong  to  the 
salt  remain  (p.  269).  The  expulsion  of  a weak  acid  from  its  salts  by  a stronger  acid 
is  similarly  completed.  While  the  neutral  salts  are  approximately  equally  dissociated, 
feeble  acids  on  the  other  hand  are  changed  but  little.  If  therefore  a strong  acid  be 
added  to  the  solution  of  such  a salt  its  hydrogen  atoms  will  come  together  with  the 
anions  of  the  salt  and  will  combine  more  or  less  completely  with  them  to  an  acid  which 
is  not  dissociated  ; there  will  remain  the  kation  of  the  salt  and  the  anion  of  the  added 
acid,  e.  g.^  hydrochloric  acid  and  sodium  borate: 

— + — + -l-  — 

Cl  + H + BO2  + Na  = HBO2  + Na  -f  Cl. 

Salts  of  weak  acids  or  of  feeble  bases  are  also  hydrolytically  decomposed  by  the  action 
of  water,  i,  e.,  they  break  down  into  base  and  acid,  of  which  the  weak  portion  exists 
undissociated  in  the  solution.  The  base  or  acid  dissociated  to  a greater  extent  may  be 
recognized  by  the  fact  that  its  solution  reacts  acid  (with  feeble  base)  or  alkaline  (with 
feeble  acid). 

The  reactions  employed  to  detect  substances  depend,  according  to  the  theory  of  elec- 
trolytic dissociation,  chiefly  upon  reactions  of  ions.  All  compounds,  for  example,  which 
in  aqueous  solution  yield  the  anion  Cl,  show  the  reaction  of  hydrochloric  acid,  so  far  as 
they  produce  a precipitate  of  silver  chloride  with  silver  nitrate.  When  chlorine  does  not 
appear  alone  as  an  ion,  but  as  a part  of  .such,  this  reaction  does  not  take  place.  The 
compound  Na.2PtClg,  the  .sodium  salt  of  hydrochlorplatinic  acid,  rich  in  chlorine,  does 
not  yield  a precipitate  of  silver  chloride  with  silver  nitrate,  because  in  arpieous  solution 
it  dis.sociates  into  the  anion  PtClg  and  the  kations  2Na.  The  color  of  the  solution  is 
al.so  materially  affected  by  the  ions.  The  following  therefore  is  accordingly  explained  : 

When  ferric  chloride  and  potassium  fluoride  meet  in  solution  in  equivalent  amounts  a 
whole  series  of  characteristics  belonging  to  ferric  chloride  solutions  disappear  : the  solu- 
tion is  colorless  ; iodine  is  not  set  free  from  potassium  iodide  even  after  the  addition  of 
an  acid  ; potassium  .sulphocyanide,  salicylic  acid  and  substances  which  otherwise  detect 
ferric  chloride  with  great  accuracy,  fail  to  show  anything.  Formerly  there  was  no  ex- 
planation for  the  deep-seated  transposition  which  had  evidently  occurred  in  the  liquid, 
but  it  may  be  found  according  to  the  theory  of  electrolytic  dissociation  in  the  fact  that 


272 


INORGANIC  CHEMISTRY. 


ferric  chloride  and  iwlassiuin  fluoride  arc  coinj)lelely  transposed  to  ferric  fluoride  anfl 
potassium  chloride  : 

FeCl.,  I 3KKI  3KC:1  1 Feld,, 

and  ferric  fluoride  is  not  dissociated.  'I'he  trivalent  iron  ions,  u|)on  which  the  reactions 
recorded  above  are  dependent,  are  no  lonj^er  present  in  the  solution  ; therefore,  the  reac- 
tion cannot  occur.  This  transposition  completes  itself  .so  fully  that  it  can  be  applied  in 
the  quantitative  determination  of  fluorides  by  working  with  ati  excess  of  ferric  chloride  and 
then  determining  the  portion,  not  converted  into  fluoride,  with  potassium  iodide,  upon 
which  ferric  fluoride  does  not  act  (Knobloch).  Similar  relations  will  be  cited  in  the 
following  pages. 


I.  GROUP  OF  THE  ALKALI  METALS. 


Potassium, 39-15 

Rubidium, 85.4 

CcTesium, 133 


Lithium, 7-03 

Sodium, 23.05 

(Ammonium,  NH^  = 18.07) 


The  metals  of  this  grotij^  are  decidedly  the  most  pronounced  in  metallo- 
basic  character,  and  this  constitutes  a visible  contrast  with  the  elements 
of  the  chlorine  group,  the  most  energetic  among  the  acid-forming  metal- 
loids. 

The  alkali  metals  in  physical  and  chemical  proj)erties  exhibit  great 
similarity.  They  oxidize  readily  in  the  air,  decompose  water  violently, 
even  in  the  cold,  with  the  formation  of  strong  basic  hydroxides,  which 
dissolve  readily  in  water  and  are  called  alkalies  (caustic  potash,  caustic 
soda)  ; hence  the  name  alkali  metal  {al  kaljiin,  Arabic,  meaning  the  ash 
of  sea  and  beach  plants,  and  the  extract  from  the  same).  They  are  not 
decomposed  by  ignition.  Their  chemical  energy  increases  with  increas- 
ing atomic  weight  (more  correctly  atomic  volume),  sodium  is  more  ener- 
getic than  lithium,  potassium  more  than  sodium,  and  rubidium  more  than 
potassium.  Caesium  has  not  been  studied  in  the  free  condition,  but,  judg- 
ing from  its  compounds,  it  possesses  a more  basic  character  than  rubidium. 
We  saw  in  other  analogous  groups  (of  chlorine,  oxygen,  phosphorus,  car- 
bon, and  similar  elements),  that  the  metalloidal,  electro-negative  character 
diminishes,  and  the  basic  increases  with  the  increasing  atomic  weight. 

The  specific  gravities  increase  simultaneously  with  the  atomic  weights; 
but  as  the  increase  of  the  latter  is  greater  than  that  of  the  former,  the 
atomic  volumes  (the  quotients  p.  252),  are  always  greater.  The  in- 

creasing fusibility  and  volatility  correspond  to  the  increase  of  the  atomic 
volumes  ; rubidium  distils  at  a red  heat,  while  lithium  volatilizes  only 
with  difficulty  : 


Li 

Na 

K 

Rb 

Cs 

Atomic  weight, 

Specific  gravity  (15°),  - - 

Atomic  volume,  .... 

J^'usion  teitiperature,  . . . 
boiling  temperature,  . . . 

7-03 

0.59 

1 1.9 

180° 

23-05 

0.97 

23-7 

96° 

742° 

39-15 

0.87 

45 

62.5° 

667° 

^5-4 

1-52 

56.1 

3^-5° 

133 

1.88 

70.7 

26.5° 

270° 

POTASSIUM. 


273 


Although  the  alkali  metals  exhibit  a great  similarity  in  their  chemical 
deportment,  we  discover  more  marked  relations  between  potassium,  rubid- 
ium and  caesium  upon  the  one  hand,  and  lithium  and  sodium  on  the 
other,  which  accords  with  their  position  in  the  periodic  system  of  the 
elements  (p.  246).  Especially  is  this  noticed  in  the  salts.  The  first 
three  metals  form  difficultly  soluble  tartrates  and  chlorplatinates  (see 
Platinum).  Their  carbonates  deliquesce  in  the  air,  while  those  of  sodium 
and  lithium  are  stable  under  similar  circumstances ; the  last  is,  indeed, 
rather  insoluble  in  water.  The  phosphates  deport  themselves  similarly; 
lithium  phosphate  is  very  difficultly  soluble.  It  must  be  remarked  that 
the  normal  carbonates  and  phosphates  of  all  other  metals  are  insoluble  in 
water.  In  lithium,  then,  which  possesses  the  lowest  atomic  weight,  it 
would  seem  the  alkaline  character  has  not  reached  its  full  expression,  and 
it  in  many  respects  approaches  the  elements  of  the  second  group,  espe- 
cially magnesium,  just  as  beryllium  approaches  aluminium.  The  elements 
of  the  two  small  periods  (lithium  and  sodium)  are,  indeed,  similar,  but 
not  completely  analogous,  while  the  homology  of  the  three  great  periods 
finds  expression  in  potassium,  rubidium  and  caesium. 


Of  the  thermo-chemical  relations  of  the  alkali  metals  only  the  heat  of  formation  of  some 
hydrates  will  be  given  here  : 

(Li,  O,  H,  Aq)  = 117.4  (Na,  O,  H,  Aq)  ^ 111.8  (K,  O,  H,  Aq)  = 116.4. 

On  comparing  these  values  with  the  heat  of  formation  of  water  ( Hg,  O)  = 68.36  Cal., 
we  immediately  perceive  why  it  is  so  readily  decomposed  by  the  alkali  metals.  All 
metals,  disengaging  more  than  68.3  Cal.  in  the  formation  of  their  oxides,  Me^O,  or  their 
hydroxides,  MeOH,  decompose  water,  and  the  energy  will  be  greater,  the  greater  the 
difference  of  heat.  The  insolubility  of  the  oxides  constitutes  an  obstacle  to  the  action  ; 
this,  however,  may  be  removed  (see  Aluminium)  by  addition  of  neutral  solvents.  Con- 
versely, all  oxides,  affording  less  heat  in  their  formation,  are  easily  reduced  by  hydrogen 
(pp.  92,  95). 


1.  POTASSIUM. 

K = 39-15- 

In  nature,  potassium  is  found  principally  in  silicates,  viz.  : feldspar 
and  mica.  By  the  disintegration  of  these  frequently  occurring  minerals, 
potassium  passes  into  the  soil,  and  is  absorbed  by  plants  ; the  ashes  of  the 
latter  consist  chiefly  of  different  potassium  salts.  The  chloride  and  sul- 
phate are  also  found  in  sea-water,  and  in  large  deposits  in  Stassfurt,  at 
Magdeburg,  and  in  Brunswick  and  in  Galicia,  where  they  were  left  by  the 
evaporation  of  the  water  of  inclosed  seas. 

Metallic  potassium  was  first  obtained  by  Davy,  in  the  year  1807,  by  the 
decomposition  of  the  hydroxide,  by  means  of  a strong  galvanic  current. 
At  present  it  is  prepared  by  igniting  an  intimate  mixture  of  carbon  and 
potassium  carbonate  : 


K.COj  T 2C  = 2K  + 3CO. 


274 


INORGANIC  CHEMISTRY. 


Such  a mixture  may  be  made  l)y  the  carl)Ouizatiou  of  organic  potassium 
salts,  <f.  , crude  tartar.  It  is  tlieu  ignited  to  wliitc  lieat,  in  an  iron 

retort,  and  tlie  escaping  vapors  collected  in  receivers  of  ]>eculiar  construc- 
tion, filled  with  rock  oil,  'I'he  latter,  a mixture  of  hydrocarbon,  serves 
as  the  best  means  of  preserving  potassium,  which  would  otherwise  oxidize 
in  the  air,  and  dccom[)ose  other  licpiids.  Potassium  carbon  monoxide 
(KC())p  is  a by-product  in  the  preparation  of  the  metal  ; see  Richter’s 
Organic  Chemistry. 

In  a fresh  section,  potassium  shows  a silver-white  color  and  brilliant 
metallic  luster.  At  ordinary  temperatures  it  is  soft,  like  wax,  and  may 
be  easily  cut.  It  crystallizes  in  octahedra,  and  has  a specific  gravity  of 
0.87  at  15°.  It  melts  at  62°  and  boils  at  about  667°,  and  when  raised  to  a 
red  heat,  is  converted  into  a greenish  vapor.  It  oxidizes  in  the  air,  and 
becomes  dull  in  color  ; heated,  it  burns  with  a violet  flame.  It  decom- 
poses water  energetically,  with  formation  of  iiotassium  hydroxide  and  the 
liberation  of  hydrogen.  If  a piece  of  the  metal  be  thrown  upon  water, 
it  will  swim  on  the  surface  with  a rotary  motion  ; so  much  heat  is  dis- 
engaged by  the  reaction  that  the  generated  hydrogen  and  the  potassium 
inflame.  Finally,  a slight  explosion  usually  results,  whereby  pieces  of 
potassium  are  tossed  here  and  there;  it  is  advisable,  therefore,  to  execute 
the  experiment  in  a tall  beaker  glass,  covered  with  a glass  plate.  Potas- 
sium combines  directly  and  very  energetically  with  the  halogens. 

On  conducting  hydrogen  over  metallic  potassium  heated  to  300-400°, 
potassium  hydride,  K2H  (or  K^H2),  results.  This  is  a metallic,  shining, 
brittle  compound,  which,  upon  stronger  heating  (above  410°),  more 
readily  in  vacno,  is  again  decomposed.  The  sodium  hydride,  Na^H2, 
obtained  in  the  same  way,  does  not  ignite  spontaneously,  but  in  other 
respects  is  very  much  like  potassium  hydride. 


The  influence  of  heat  and  pressure  in  the  formation  and  decomposition  of  these  com- 
pounds is  very  noteworthy.  If,  for  example,  potassium  hydride  be  heated  it  melts,  but 
otherwise  remains  unchanged.  Above  2DO°  (in  a vacuum)  it  sustains  a partial  decom- 
position (dissociation),  which  gradually  increases  as  the  temperature  rises.  If  the  heating 
should  take  place  in  a closed  vessel  provided  with  a manometer,  it  will  be  observed  that 
the  decomposition  at  a given  temperature  will  continue  until  the  liberated  hydrogen  has 
acquired  a definite  tension — until  it  exerts  a definite  pressure.  For  potassium  hydride, 
this  tension,  at  330°,  equals  45  mm.  The  decomposition  will  then  cease,  but  will  pro- 
ceed further  at  the  same  temperature  if  the  hydrogen  gas  be  removed,  until  the  pressure 
of  45  mm.  is  again  reached.  This  pressure  is  therefore  called  the  dissociation  tension. 
In  this  manner  a complete  decomposition  of  the  hydride  may  be  effected  at  the  tempera- 
ture given  above.  If,  however,  the  di.sengaged  hydrogen  is  not  removed,  but  be  added 
to  the  completely  or  jrartially  decomjrosed  hydride,  and  the  pressure  be  raised  to  45  mm. 
(at  the  tem])erature  330°),  the  jiotassium  hydride  will  be  re-formed.  Consequently,  both 
the  decomposition  and  the  formation  of  a body  can  follow,  dejrending  uj)on  whether  the 
external  partial  ])ressure  be  lowered  or  increased.  Similar  phenomena  occur  at  higher 
temperatures,  tlie  corres|)ondiiig  pressure,  of  course,  increasing  by  regular  steps. 

'I'he  tension  of  dissociation  is  independent  of  the  relative  (juantity  of  the  dissociated 
body  and  of  the  s[)ace  which  the  di.sengaged  gas  can  occupy,  whereas  in  solutions  and 
absori)li()ns  (ammonia  by  charcoal)  the  pressure  at  one  and  the  same  temperature 
increases  with  (he  (|uantity  of  the  absorbed  gas. 

All  exothermic  compounds  behave  like  potassium  and  .sodium  hydrides  when  they  are 
decomposed  into  their  eoin|)()nents  ; if  the  pressure  be  raised  above  the  tension  of  disso- 
ciation the  components  reunite — the  compounds  are  re-formed.  The  decomposition  of 


POTASSIUM  HYDROXIDE. 


275 


the  endothermic  compounds  (potassium  chlorate  into  chloride  and  oxygen)  is  quite  dif- 
ferent (pp.  30,  94).  It  proceeds  with  heat  disengagement  (KCljO.^  = — '9.8  Cal.), 
corresponding  to  the  chemical  affinities,  and  is  only  induced  by  application  of  external 
heat.  It  is  independent  of  external  pressure,  and  there  is  no  reunion  of  the  decomposi- 
tion products  upon  increasing  the  external  pressure  or  upon  lowering  the  temperature. 

We  must  not  omit  mentioning  the  great  analogy  between  the  phenomena  of  dis- 
sociation and  the  vaporizing  of  liquids.  Like  dissociated  bodies,  liquids  exhibit  at  all 
temperatures  a definite  vapor  tension.  If  the  pressure  above  the  liquid  be  diminished 
the  evaporation  will  continue  until  the  vapor  tension  is  regained,  but  if  the  pressure  be 
increased  then  a corresponding  portion  of  the  vapor  will  be  condensed. 


Oxygen  compounds  of  potassium  are  not  known  in  a pure  state.  The 
metal  is  not  attacked  by  pure,  dry  oxygen  below  60-80°.  Above  this 
temperature  it  burns  in  the  gas,  if  its  surface  be  renewed,  to  a yellow 
mass  which,  according  to  Erdmann  and  Kothner,  consists  of  the  sesqui- 
oxide,  K2O3,  and  superoxide,  KO2.  The  oxide  K2O,  from  which  the 
potassium  salts  are  derived,  is  not  definitely  known. 

Potassium  Hydroxide,  o,r  Caustic  potash  {^Kalium  causticuni), 
KOH,  is  obtained  by  the  action  of  potassium  or  its  oxide  upon  water.  For 
its  preparation,  potassium  carbonate  is  decomposed  by  calcium  hydroxide 
(slaked  lime)  ; 

K2CO3  + Ca(OH)2  = CaC03  + 2KOH. 

The  solution  of  i part  of  potassium  carbonate  in  10-12  parts  of  water  is  boiled  with  i 
part  of  slaked  lime  in  an  iron  pot  until  a filtered  portion  does  not  effervesce,  when  hydro- 
chloric acid  is  added  ; i.  e.,  until  there  is  no  longer  any  carbonic  acid  present.  On 
standing  awhile,  the  insoluble  calcium  carbonate  subsides,  and  the  liquid  becomes  clear. 
The  solution  of  potassium  hydroxide  is  then  poured  off,  evaporated,  the  residue  melted  in 
a silver  dish  (which  it  does  not  attack),  until  the  hydroxide  begins  to  volatilize  in  clouds, 
when  it  is  poured  into  moulds  [K.  causticuni  ficsum).  The  caustic  potash,  prepared  in 
this  way,  is  not  entirely  pure,  but  contains  potassium  chloride  and  other  salts.  To 
obtain  a product  that  is  chemically  pure,  fuse  potassium  nitrate  with  copper  filings,  and 
evaporate  the  aqueous  extract  of  the  fusion  in  silver  vessels. 

One  of  the  most  brilliant  achievements  of  modern  chemical  industry  is  the  electrolytic 
decomposition  of  the  chlorides  of  the  alkali  metals  into  free  chlorine  and  7uetal ; the  latter  at 
the  moment  of  its  liberation  is  converted  by  water  into  its  hydroxide.  The  practical  solu- 
tion of  this  problem  has  made  this  operation  a technical  process  since  1890.  The  Stass- 
furt  potassium  chloride  has  been  thus  made  to  yield  pure  caustic  in  solution  and  in  solid 
form  ; also  chlorine,  bleaching  lime  and  hydrogen  (pp.  51,  278  ; also  under  Soda). 

Potassium  hydroxide  forms  a white,  crystalline  mass  which  fuses  rather 
easily,  and  volatilizes  undecomposed  at  a very  high  temperature.  Ex- 
posed to  the  air  it  deliquesces,  as  it  absorbs  water  and  carbon  dioxide. 
It  is  very  soluble  in  alcohol,  and  esjiecially  in  water.  The  solution 
{Liquor  kalii  caustici)  possesses  a strong  alkaline  reaction,  saponifies  the 
fats,  and  has  a corrosive  action  upon  the  skin  and  organic  tissues.  At 
low  temperatures  the  hydrate  KOH  -f-  2H2O  crystallizes  out  from  con- 
centrated solutions. 


The  haloid  salts  of  potassium  are  obtained  by  the  direct  union  of  the 
halogens  with  potassium,  and  by  the  saturation  of  the  hydroxide  or  car- 


276 


INORGANIC  CHEMISTRY. 


bonatc  witli  lialoid  acids,  d'hey  arc  readily  solulile  in  water,  have  a 
salty  taste,  and  crystallize  in  cul)es.  When  heated  tliey  melt,  and  are 
somewhat  volatile. 

Potassium  Chloride  (A".  chloratuj?i),  K.C1,  occurs  in  Stassfurt  in 
large  deposits,  as  sylvite,  and  c(;mbined  with  magnesium  chloride  exists 
di?,  carnaili^e  {iAgC\^,  KCl  -f  6H,/)). 

The  opening  up  and  the  practical  yields  of  the  Stassfurt  salt  beds  containing  carnal- 
lite  ; kieserite,  MgS(  )^ . 1 1.^( ) ; tachhydrite,  CaCl.^ . 2MgCl2  + 12II./);  kainite,  MgSO^.- 
K.^SO^.  MgCl.^  hll.^O;  boracite,  2Mg3H^r)|.  j b?  ^ii^l  ^1''^’  brotnides  (p.  53)  have 
been  of  great  coinniercial  importance  to  the  (ierinan  1‘hnpire.  'These  arc  the  largest  salt- 
beds  in  the  world.  before  they  were  worked  very  considerable  (|uantities  of  salt  were 
imported  into  Germany  but  now  that  country  stands  at  the  head  of  all  the  salt  producing 
portions  of  the  earth.  In  1897  .Stassfurt  yielded  273,364  tons  of  rock-salt  ($277,026) 
and  1,946,188  tons  of  crude  jwtash  salts  ($6,513,553)  ; in  the  same  year  the  production 
of  potassium  chloride  in  the  German  Empire  amounted  to  108,000  tons  ($5,764,423). 

Carnallite  serves  as  the  chief  source  for  the  pre|)aration  of  potassium 
chloride;  water  decomposes  it  into  the  more  sparingly  soluble  jxjtassium 
chloride  and  the  readily  soluble  magnesium  chloride.  It  is  interesting  to 
note  that  three-fourths  of  the  ])Otassium*  chloride  sejtarate  in  solid  form 
when  carnallite  is  heated  to  176°.  The  licpiid  separated  from  this  yields 
carnallite  again  on  cooling  to  115°,  while  magnesium  chloride  remains 
dissolved.  The  chloride  crystallizes  in  vitreous  cubes,  of  sj^ecific  gravity 
1.98.  It  melts  at  about  800°,  and  volatilizes  at  a strong  red  heat  [see 
Jahrbuch  der  Chemie  v (1895),  66,  67].  100  parts  of  water  dissolve  29 

parts  of  the  salt  at  0°,  and  56  parts  at  100°.  Potassium  chloride  is  used 
in  making  nearly  all  the  other  potassium  salts,  hence  it  is  largely  ap- 
plied technically  (see  Potassium  Carbonate,  Potassium  Nitrate,  Potassium 
Chlorate). 

Potassium  Bromide  (^K.  bromatuni),  KBr,  is  generally  obtained 
by  warming  a solution  of  potassium  hydroxide  with  bromine,  when  the 
bromate  is  also  produced  : 

6KOH  3Br.2  = 5KBr  -|-  KBrOg  y 3H2O. 

The  solution  is  evaporated  to  dryness,  mixed  with  charcoal,  and  ignited, 
which  reduces  the  bromate  to  bromide  : 

KBrOg  -)-  3C  = KBr  -|-  3CO. 

It  is  readily  soluble  in  water  and  slightly  in  alcohol ; forms  cubes  of  spe- 
cific gravity  2.4,  and  melts  at  about  740°. 

Potassium  Iodide  (A',  iodatuni),  KI,  may  be  prepared  like  the 
})receding.  The  iodate  produced  along  with  the  iodide  upon  adding 
iodine  to  caustic  potash  may  be  reduced  by  hydrogen  peroxide  : 

KlOg  + 3IIA  = d 3II2O  + 3O,  (p.  180). 

It  is  usually  obtained  according  to  the  following  method  : Iodine  (3 
parts)  and  iron  filings  (i  part)  are  rubbed  together  under  water  ; an  equal 
(|uantity  of  iodine  (i  part)  is  again  added  to  the  solution  of  this  ferrous 
iodide,  b'el^,  in  order  to  form  ferrous-ferric  iodide,  Fe.Jg,  which  is  then 
boiled  with  the  recpiired  (juantity  of  pota.ssium  carbonate;  this  will  pre- 


POTASSIUM  CHLORATE. 


277 


cipitate  ferrous-ferric  oxide ; carbon  dioxide  escapes,  and  potassium 
iodide  will  be  found  in  the  solution  ; 

Fcglg  -f  4K2CO3  = FegO^  + SKI  + 4CO2. 

It  forms  large  white  crystals,  fuses  at  about  720°,  and  is  tolerably  volatile. 
Its  specific  gravity  equals  3.0.  At  medium  temperatures  it  dissolves  in 
0.7  part  of  water  and  2.5  parts  of  absolute  alcohol.  The  aqueous  solution 
dissolves  iodine  in  large  quantity.  Many  metallic,  insoluble  iodides  dis- 
solve in  it  without  difficulty,  forming  double  iodides,  e.  g.,  Hgl2.2KI. 
The  iodide  is  employed  in  medicine  and  in  photography. 

Potassium  Fluoride,  KFl,  is  obtained  by  dissolving  the  carbonate  in  aqu^us  hydro- 
fluoric acid.  It  crystallizes  in  cubes  at  ordinary  temperatures,  with  two  molecules  of 
water,  but  above  35°  does  not  contain  water  of  crystallization.  It  is  very  .soluble  in 
water.  The  aqueous  solution  attacks  glass.  It  is  greatly  inclined  to  combine  with  other 
fluorides:  KFl.HFl;  BFI3.KFI.  The  commercial  salt  is  frequently  rich  in  arsenic. 
On  adding  hydrofluosilicic  acid  to  the  solution  of  potassium  .salts,  a gelatinous  precipitate 
of  potassium  silico-fluoride,  K2SiFlg,  is  thrown  down,  wdiich  dissolves  with  difficulty  in 
water. 

Potassium  Cyanide,  KCN.  This  salt  can  be  produced  by  saturating  potassium 
hydroxide  with  hydrocyanic  acid,  and  by  heating  yellow  prussiate  of  potash  (see  Iron). 
It  forms  a white,  easily  fu.sible  mass,  which  deliquesces  in  the  air.  The  solution  may  be 
easily  decomposed.  The  salt  crystallizes  in  cubes,  has  an  alkaline  reaction,  and  smells 
like  prussic  acid.  As  the  result  of  hydrolysis  free  prussic  acid — which  is  very  slightly 
dissociated — is  pre.sent  with  free  alkali.  The  introduction  of  carbon  dioxide  completes  the 
decomposition.  By  fusion  potassium  cyanide  reduces  many  metallic  oxides,  and  hence  is 
employed  in  reduction  proces.ses.  It  is  just  as  poisonous  as  prussic  acid.  It  is  applied 
in  many  ways,  especially  in  photography  and  for  galvanic  silvering  and  gilding.  Lately 
it  has  met  with  extended  application  in  the  extraction  of  gold  from  low-grade  ores  and 
from  sand.  Generally  it  is  a mixture  of  potassium  and  sodium  cyanides  which  is  used  for 
this  purpose  ; this  can  be  obtained  by  fusing  potassium  ferrocyanide  with  sodium  : 

K,Fe(CN)6  -f  2Na  = 2NaCN  + 4KCN  -f  Fe. 

Potassium  Chlorate  (A",  chloricuni)  KCIO3,  is  produced  when  a 
slight  excess  of  chlorine  is  introduced  into  caustic  potash,  and  the  hypo- 
chlorite, formed  at  first,  oxidized  thereby  to  chlorate.  This  change  pro- 
ceeds most  rapidly  at  80-90°  (see  p.  178).  The  sparingly  soluble  chlorate 
separates  when  the  solution  cools. 

Technically,  a solution  of  calcium  chlorate,  produced  by  slightly  super- 
saturating lime-water  with  chlorine  at  40°,  is  mixed  with  a sufficient  quan- 
tity of  ])otassium  chloride,  when  potassium  chlorate  and  calcium  chloride 
result : 

Ca(C103)2  + 2KCI  = CaCb  + 2KCIO3. 

Magnesia  can  be  advantageously  substituted  for  lime-water. 

At  present  this  old  method  of  Liebig  is  being  more  and  more  supplanted 
by  the  electrolytic  method,  particularly  since  Oettel  found  that  potassium 
chlorate  is  formed  electrolytically  in  alkaline  solutions  of  ])otassium  chlo- 
ride without  the  use  of  a diaphragm,  which  se])arated  the  anode  from  the 
kathode  liquor.  Since  the  electrolysis  of  potassium  chloride  yields  caustic 
(together  with  hydrogen)  at  the  kathode  and  chlorine  at  the  anode,  the 
conditions  for  the  production  of  chlorate  are  evident.  See  also  Elbs, 
Chem.  Zeit.  1897,  996. 


INORGANIC  CHEMISTRY. 


278 


A French  company  in  Switzerland  first  made  cldorate  electrolytically,  while  the  caustic 
alkalies  were  first  produced  in  this  way  in  (Jermany;  England  followed  later.  'I’he  ledi 
nical  manufacture  of  electrolytic  chlorate  has  been  in  operation  since  1891  in  Switzerland. 
The  chemical  factory  at  (Iriesheiin  began  to  make  caustic  potash  and  soda  electrolytically 
in  1890,  Since  1894  the  electric  factory  at  Hitterfeld  has  jmxluced  chiefly  caustic  and 
bleaching  li(juors  in  the  electrolytic  way.  Other  companies  have  followed  those  men- 
tioned, and  by  some  of  them  soda  is  now  being  prepared  from  electrolytic  caustic  soda. 


Potassium  chlorate  crystallizes  from  the  hot  solution  in  shinitig  tables 
of  the  monoclinic  system,  which  dissolve  with  difficulty  in  cold  water(Too 
parts  at  the  ordinary  temperature  dissolve  6 ])arts  of  the  salt).  Its  taste 
is  cooling  and  astringent.  When  heated  it  melts  at  360°  and  at  higher 
temperatures  gives  up  a ])ortion  of  its  oxygen,  and  changes  to  the  chloride 
and  Perchlorate,  KCK)^,  which  on  further  heating  decomposes  into 
oxygen  and  potassium  chloride  (see  \).  178).  With  hydrochloric  acid  it 
liberates  chlorine : 

KCIO3  + 6IIC1  = KCl  + 3H2O  + 3CI2. 

Mixed  with  sulphur,  or  certain  sulphides,  it  explodes  on  heating  and 
when  struck  a sharp  blow.  The  igniting  material  upon  the  so-called 
Swedish  (parlor)  matches  consists  of  antimony  sulphide  and  potassium 
chlorate;  when  this  is  rubbed  upon  the  friction  surface  coated  with  red 
phosphorus  it  ignites. 

Potassium  Hypochlorite,  KCIO,  is  formed  when  chlorine  is  con- 
ducted into  a solution  of  potassium  hydroxide : 

2KOH  + CI2  = KCl  + KCIO  -f  H2O. 

It  only  exists  in  aqueous  solution;  when  the  latter  is  evaporated  the 
salt  is  decomposed  into  chloride  and  chlorate : 

3KCIO  = 2KCI  4-  KCIO3. 

In  the  presence  of  an  excess  of  chlorine  chlorate  is  rapidly  produced. 

The  solution  has  an  odor  resembling  that  of  chlorine,  and  bleaches 
strongly,  especially  upon  the  addition  of  acids  (p.  175).  The  bleaching 
solutions  occurring  in  trade  (Eau  de  Javelle)  are  prepared  by  the  action  of 
chlorine  upon  solutions  of  sodium  (Eau  de  Labarraque)  and  potassium 
(Eau  de  Javelle)  carbonates;  but  recently  they  have  been  made  by  the 
electrolysis  of  the  corresponding  chlorides.  They  also  contain  free' hypo- 
chlorous  acid. 

The  oxy-salts  of  bromine  and  iodine  are  perfectly  analogous  to  those  of  chlorine. 
Potassium  hromnte,  KBrO,,  and  Potassium  iodate,  KIO3,  are  ])repared  by  the  action  of 
bromine  or  iodine  u])on  hot  potassium  hydroxide  ; the  second  is  also  produced  by  the 
action  of  iodine  upon  ])otassium  chlorate,  when  the  chlorine  is  directly  replaced  (p.  iSol. 
'The  exj)erimcnts  of  Klinger  and  of  Bassett  apparently  prove  that  a direct  replacement  of 
chlorine  by  iodine  does  not  take  place  here  ; it  is  rather  the  oxidation  of  the  latter  by  the 
anion  (.'Kqof  the  chlorate.  I^)tassiuln  iodate  can  also  be  made  by  oxidizing  potassium 
iodide  (i  |)art)  with  ])otassium  ])ennanganate  (2  parts)  in  aqueous  solution.  If  chlorine 
be  passed  through  a hot  solution  of  jiotassium  iodate  or  iodide  in  jK)tassium  hydroxide, 
the  p(;rioda((‘  of  potassium,  KIO,,  arises;  it  is  difficultly  soluble  and  when  heated  de- 
composes into  oxygen  and  potassium  iodate,  which  then  breaks  down  into  potassium 
iodide  and  oxygen. 


POTASSIUM  NITRATE. 


279 


Besides  the  normal  periodates,  KIO^,  NalO^,  other  salts  exist  which  are  derived  from 
the  highest  hydroxyl  compound,  and  its  anhydro- derivatives  (p.  18 1).  These 

salts  are  very  numerous,  and  are  in  part  monoperiodates,  10(011)5  and  102(011)3,  and 
partly  polyperiodates,  produced  by  the  condensation  of  several  molecules  of  the  highest 
hydroxides  with  the  exit  of  water,  e.  g.^  l203(0H)g,  1205(011  )4,  and  l20g(0Il)2. 

Potassium  Sulphate,  K2SO4,  is  formed  in  the  action  of  concentrated 
sulphuric  acid  upon  potassium  chloride,  and  as  a by-product  in  many 
technical  operations.  It  is  also  obtained  by  transposing  schonite, 
MgSO^.  K2SO4 -f  6H2O,  and  other  Stassfurt  salts  containing  sulphates 
with  potassium  chloride: 

MgSO,.K2SO,  -f  2KCI  = 2K2SO,  -f  MgCb- 

It  crystallizes  without  water,  in  small  rhombic  prisms,  having  a bitter, 
salty  taste,  and  dissolves  in  10  parts  of  water  at  the  ordinary  temperature. 
It  melts  at  about  1070°.  It  is  employed  principally  for  the  preparation 
of  potassium  carbonate,  according  to  the  method  of  Le  Blanc  (p.  281). 

The  acid  or  primary  salt,  KHSO^,  crystallizes  in  large  rhombic 
tables,  and  is  very  readily  soluble  in  water.  It  fuses  at  about  200°,  loses 
water,  and  is  converted  into  potassium  pyrosulphate,  K2S20^,  which  at 
600°  yields  K2SO4  and  SO3  (p.  194). 

Potassium  Sulphite. — The  salts  of  sulphurous  acid— the  primary^  KHSO3,  and 
the  secondary  sulphites,  K2SO3 — are  produced  when  sulphuric  dioxide  comes  in  contact 
with  a potassium  carbonate  solution  ; they  are  very  soluble  and  crystallize  with  difficulty. 
The  first  salt  shows  an  acid,  the  second  an  alkaline  reaction.  If  sulphur  dioxide  be 
passed  into  a solution  of  potassium  carbonate  until  effervescence  ceases  and  then  cooled, 
the  pyrosulphite,  K2S.2O5,  corresponding  to  the  pyrosulphate,  will  crystallize  out. 

Potassium  Persulphate,  K2S20g,  results  on  electrolyzing  a saturated  solution  of 
acid  potassium  sulphate,  when  it  separates  at  the  anode  as  a white  crystalline  precipitate. 
It  can  be  crystallized  from  hot  water  ; on  rapid  cooling  it  separates  in  small  prisms.  Its 
solution  oxidizes,  has  a cooling,  salt-like  taste  and  does  not  yield  a precipitate  with  solu- 
tions of  other  metals  (the  salts  of  silver,  manganese  and  cobalt  excepted).  The  dry  salt 
commences  to  decompose  at  100°  into  oxygen  and  pyrosulphate  (see  p.  188). 

Potassium  Nitrate,  Saltpeter  {K.  nitrictmi),  KNO3,  occurs  in  the 
largest  amounts  in  East  India,  especially  in  Ceylon,  at  Bengal  and  at 
Gutscharat  (Bombay)  in  distinct  layers;  in  cavities  (in  Ceylon)  once 
the  lairs  of  animals  and  of  men’s  dwelling-places,  which  are  even  now 
inhabited  by  hosts  of  field  mice.  It  is  found  in  abundance  in  other 
torrid  regions  in  the  soil,  upon  which  it  effloresces  during  the  dry  sea- 
son (hence  sal,  salt;  iz^rpa,  rock),  e.  g.,  in  Peru,  in  Bolivia,  in  South 
Africa,  also  in  Egypt,  etc.  It  is  produced  wherever  nitrogenous  organic 
substances  decay  in  the  presence  of  potassium  carbonate  (aided  by  micro- 
organisms), conditions  which  are  present  in  almost  every  soil.  The 
intentional  introduction  of  these  is  the  basis  of  the  artificial  niter  pro- 
duction in  the  so-called  saltpeter  plantations,  which  were  formerly  cul- 
tivated actively  in  Spain,  Hungary,  Sweden  and  Switzerland.  Manures 
and  various  animal  offals  are  mixed  with  wood  ashes  (i)otassium  car- 
bonate) and  lime,  arranged  in  porous  layers,  and  submitted  to  the  action 
of  the  air  (protected  from  rain)  for  two  or  three  years,  when  nitrates  are 
produced  from  the  slow  oxidation  of  the  nitrogen.  The  heaps  are  then 


28o 


INORGANIC  CUKMISTRY. 


treated  with  water  and  potassium  carbonate  added  to  the  solution,  which 
contains  potassium,  calcium  and  magnesium  nitrates,  to  convert  the  last 
two  salts  into  potassium  nitrate  : 

Ca(N03)2  + 1^2^03  = CaCOa  + 2KNO3. 

The  precipitate  of  calcium  and  magnesium  carbonates  is  filtered  off  and 
the  solution  evaporated.  'Fhe  soils  of  J^^ast  India  containing  niter  are 
similarly  worked.  Until  the  Crimean  War  (1852-1855)  the  demand 
for  poiassium  nitrate  in  manufacturing  gunjiowder  was  met  almost  exclu- 
sively by  h^ast  India.  The  numerous  and  constant  inquiries  for  the  salt 
led  (jerman  chemists  to  transiiose  the  Chili  saljieter  by  means  of  Stass- 
furt  potassium  chloride  into  the  j)>;tassium  salt  (conversion  saltpeter): 

NaN()3  + KCl  = NaCl  | KNf)3. 

Warm  saturated  solutions  of  e(iui valent  quantities  of  sodium  nitrate 
and  potassium  chloride  are  mixed  and  boiled,  when  sodium  chloride, 
being  less  soluble  in  hot  water,  will  sejiarate.  On  cooling  the  solution 
potassium  nitrate,  being  less  soluble  in  cold  water,  crystallizes  out; 
sodium  chloride  is  about  equally  soluble  in  hot  and  cold  water,  for  which 
reason  the  portion  not  separated  by  boiling  remains  in  solution  (i).  266). 

Potassium  saltpeter  crystallizes  without  water  of  crystallization  in  large, 
six-sided  rhombic  prisms.  It  is  far  more  soluble  in  hot  than  in  cold 
water;  100  parts  of  w'ater  dissolve  247  parts  at  100°,  but  at  0°  only  13 
])arts.  It  possesses  a cooling  taste,  fuses  at  340°,  and  decomposes,  when 
further  heated,  into  oxygen  and  potassium  nitrite,  KNO2.  Heated  with 
carbon  it  yields  potassium  carbonate : 

4KNO3  -|-  5C  = 2K2CO3  -(-  3CO2  T 2N2. 

Its'principal  use  is  in  the  manufacture  of  gunpowder.  This  is  a granular  mixture  of 
potassium  nitrate,  sulphur,  and  charcoal.  The  relative  quantities  of  these  constituents 
are  somewhat  different  in  the  various  kinds  of  powder  (sporting,  blasting,  and  powder 
free  from  sulphur).  The  first  consists  of  4KNO3  -|-  2C  -f-  2S  = 2K^S0^  -[-  2CO2  + 2N2 ; 
the  second  : 4KNO3  6C  -)-  4S  — 2K2S2  -f-  6CO2  -(  2N2»  ^he  third  : 4KNO3  -j-  5C 
= 2K2CO3  -f-  3Ct)2  + 2N2.  These  three  varieties,  mixed  in  different  but  simple  propor- 
tions, constitute  the  powders  in  general  use.  Each  variety  possesses  peculiarities  as  to 
ignition,  combustibility,  energy  content,  heat  and  gas  content.  The  mixing  is  made  in 
accordance  with  the  demand  for  any  one  or  more  of  these  properties.  The  effectiveness 
of  the  powder,  therefore,  depends  upon  the  disengagement  of  carbon  dioxide  and  nitrogen 
gas,  the  volume  of  which  is  almost  1000  times  as  great  as  that  of  the  decomposed  powder. 

Potassium  Nitrite  (A".  nUrosuvi),  KNO2,  is  obtained  by  fusing 
saltpeter  with  lead  (2  parts)  which  withdraws  one  atom  of  oxygen  from 
the  former  (p.  205).  It  forms  a white,  crystalline,  fusible  mass,  which 
deliquesces  in  the  air. 

Potassium  Phosphates. — The  potassium  salts  of  phosphoric  acid : 
K^PO^,  K2lIP()^,  and  KH2P(),j,  meet  with  no  ju'actical  ai)plication,  they 
are  readily  soluble  in  water  and  crystallize  ])oorly  ; therefore,  the  sodium 
salts  are  generally  used.  The  borates,  KIIO2  and  K2B^O.  -|-  5H2O  (see 
borax),  crystallize  with  difficulty. 

Potassium  Carbonate,  K2(X).,,  ordinarily  known  as  potashes,  is 
the  principal  ingredient  of  the  ashes  of  land  plants. 


POTASSIUM  CARBONATE. 


281 


As  1000  parts  of  wood  yield  3.5-28  parts  of  ashes  and  the  potassium 
carbonate  in  the  latter  is  0.45-4  parts  it  is  obvious  that  only  countries 
like  Russia,  Canada,  the  United  States,  Hungary  and  Galicia  where 
there  is  an  abundance  of  woodland  can  produce  potashes  in  large 
amounts  from  the  wood  ashes.  The  latter  are  extracted  with  water,  the 
clear  filtrate  is  then  evaporated  until  it  solidifies  on  cooling,  when  the 
residue  is  dehydrated  and  decolorized  by  calcination  in  ovens.  The  prod- 
uct is  crude  potashes.  By  repeating  the  preceding  treatment  purified 
potashes  or  pearl  ash  is  obtained.  In  countries  like  Belgium,  Germany, 
France  and  Switzerland  where  sugar  beets  are  cultivated  large  quantities 
of  potassium  carbonate  are  separated  from  the  ash  of  the  beets.  Beets 
withdraw  large  quantities  of  potash  salts  from  tl.e  soil  which  if  the  latter 
is  to  remain  fertile  must  be  returned  to  it  by  potash  fertilizers.  In  this 
direction  the  Stassfurt  potash  salts  are  most  valuable.  The  alkali  salts  of 
the  beet  are  in  its  juice  and  they  remain,  when  the  sugar  is  extracted 
from  the  latter,  in  the  molasses.  When  the  latter  is  allowed  to  ferment, 
then  evaporated  and  subjected  to  dry  distillation,  alcohol,  ammonia, 
trimethylamine  and  other  valuable  substances  are  obtained  ; furthermore, 
the  residual  coke  is  rich  in  potash.  Its  aqueous  extract  is  worked  for 
potassium  carbonate.  Sheep’s  suint  contains  potassium  salts  of  organic 
acids  which  on  incineration  yield  potassium  carbonate.  The  same  salt 
is  obtained  as  a by-product  in  the  manufacture  of  iodine  and  bromine 
from  sea-algse  and  sea-weeds. 

In  Germany  all  these  methods  for  the  preparation  of  potassium  carbon- 
ate give  place  to  those  in  which  the  Stassfurt  potassium  chloride  is  util- 
ized, from  which  potassium  carbonate  is  prepared  by  two  methods  : 

1.  Method  of  Le  Blanc. — This  will  be  discussed  in  connection  with 
soda.  H.  Griineberg  was  the  first  to  employ  it  in  making  potashes. 

2.  Method  of  Ch.  R.  Engel  in  Montpellier. — Brecht  introduced  this 
method  into  German  industries.  Magnesium  carbonate  is  mixed  with  a 
solution  of  potassium  chloride,  and  the  liquid,  while  being  stirred,  is 
saturated  with  carbon  dioxide.  A double  salt  of  magnesium  carbonate 
and  acid  potassium  carbonate  separates  while  the  solution  contains  mag- 
nesium chloride : 

SMgCOj  + 2KCI  + CO2  + 9H2O  = 2[MgC03 . KHCO3. 4H2O]  + MgCh- 

The  double  salt  is  freed  from  any  adherent  liquor  by  washing  it  with  a 
solution  of  magnesium  bicarbonate,  after  which  it  is  decomposed  under 
pressure  with  water  at  140°.  Basic  magnesium  carbonate  separates  in  a 
compact  form ; carbon  dioxide  escapes,  and  the  liquid  containing  pure 
potashes  is  evai)orated  and  the  residue  calcined.  Potassium  carbonate 
free  from  sodium  salts  can  be  obtained  in  this  way,  because  sodium  chlo- 
ride does  not  act  upon  magnesium  carbonate. 

The  commercial  carbonate  is  a white,  granular,  deliquescent  powder, 
melting  at  890°,  and  vaporizing  at  a red  heat.  It  crystallizes  from  con- 
centrated aqueous  solutions  with  molecules  of  water,  in  monoclinic 
prisms;  at  100°  it  loses  ^ molecule  of  water.  The  solution  has  a caustic 
taste  and  shows  an  alkaline  reaction.  When  carbon  dioxide  is  con- 
24 


282 


INORGANIC  CHEMISTRY. 


ducted  through  the  liquid  it  is  absorl)cd  and  primary  potassium  carbonate 
is  produced  : 

KjCO, -f  1 1/)  I CO,  = 2KIICO,. 

This  salt,  ordinarily  called  bi-carbonate,  crystallizes  in  inonoclinic 
l)risms,  free  from  water.  It  dissolves  in  3-4  ])arts  of  water  and  exhibits  a 
neutral  reaction.  When  heated,  it  decomi)oses  into  i)otassiuin  carbonate, 
carbon  dioxide,  and  water,  'bhe  decornj)osition  of  the  dry  salt  does  not 
begin  until  at  about  110°,  while  the  aqueous  solution  decomposes  even  on 
evaporation.  Potassium  carbonate  is  used  almost  entirely  in  the  produc- 
tion of  Bohemian  or  crystal  glass. 

Constam  and  v.  Hansen  electrolyzed  solutions  of  potassium  carbonate,  cooled  to 
— 15°,  and  obtained  at  the  anode  Potassium  Percarbonate,  K2C20g, — a slightly  blue- 
colored,  deliquescent  powder.  Jt  is  formed,  like  the  persulpliates,  by  the  union  of  the 
ions  KCO.J,  into  which  (together  with  potassium  ions)  the  carbonate  is  decomposed  by 
the  current.  It  resembles  the  persuljdiates  in  properties,  in  so  far  that  when  heated  to 
200-300°  it  rapidly  decomposes  into  carbonate  and  oxygen.  Oxygen  escapes  from  its 
aqueous  solution  at  45°.  It  is  a powerful  oxidant  ; many  dyes  are  bleached  by  it. 
Dilute  acids  evolve  hydrogen  peroxide  from  it:  K2C2f)fi  4 2iICl  2CO2  + 2KCI -|- 
H2O2 ; this  also  occurs  with  caustic  potash  : K2C20g  | 2KOII  = 2K2CO^  -j-  H/),.  Its 
solution  rapidly  reduces  manganese  peroxide,  lead  peroxide  and  silver  oxide  with  energetic 
liberation  of  oxygen  : Ag20  f K2C20g  = Ag,  -f-  K2CO3  4-  CO,  -(-  O,.  See  p.  102.  Its 
chemical  structure  is  very  likely  analogous  to  that  of  the  persulphates  (pp.  i88,  189,  279)  : 

O — SO3K  O — CO3K 

6 — SO3K  6 — CO3K. 

Potassium  Silicate,  water-glass,  does  not  possess  a constant  com- 
position and  cannot  be  obtained  crystallized.  It  is  produced  by  solution 
of  silicic  acid  or  amorphous  silicon  dioxide  in  potassium  hydroxide,  or  by 
the  fusion  of  silica  with  potassium  hydroxide  or  carbonate.  The  concen- 
trated solution  dries  in  the  air  to  a glassy,  afterward  opaque,  mass, 
which,  when  reduced  to  a powder,  will  dissolve  in  boiling  water.  Potas- 
sium (and  also  sodium)  water-glass  has  an  extended  application,  especially 
in  cotton  printing,  for  the  fixing  of  colors  (stereochromy),  in  rendering 
combustible  material  fireproof,  in  soap  boiling,  etc. 


SULPHUR  COMPOUNDS  OF  POTASSIUM. 

Potassium  Hydrosulphide,  KSH,  is  obtained  when  potassium 
hydroxide  is  saturated  with  hydrogen  sulphide  : 

KOH  + II,S  = KSH  -b  H2O. 

Evaporated  in  vacuo  it  crystallizes  in  colorless  rhombohedra,  of  the  for- 
mula 2KSri  -j-  H,0,  which  deliquesce  in  the  air.  At  200°,  it  loses  its 
water  of  crystallization,  and  at  a higher  temperature  fuses  to  a yellowish 
li(]uid,  which  solidifies  to  a reddish  mass.  Like  the  hydroxide,  it  has  an 
alkaline  reaction.  On  adding  an  ecpiivalent  quantity  of  potassium  hydrox- 
ide to  the  sulphydrate  solution,  we  get  potassium  sulphide: 

KSH  + KOH  = K,S  -f  H,,0. 


SULPHUR  COMPOUNDS  OF  POTASSIUM. 


283 


Potassium  Sulphide,  KgS,  is  usually  obtained  as  a porous  mass  by 
gently  heating  a mixture  of  potassium  sulphate  and  carbon  in  well-closed 
crucibles : 

KjSO^  + 2C  = KjS  + 2CO2. 

When  fused,  it  solidifies  to  a red  crystalline  mass.  It  crystallizes  from 
concentrated  aqueous  solutions  with  five  molecules  of  water,  in  colorless 
prisms,  which  deliquesce  in  the  air.  The  solution  absorbs  oxygen  from 
the  air,  and  is  decomposed  into  potassium  hyposulphite  and  caustic 
potash : 

2KjS  -h  HjO  + 2O2  = K2S2O3  + 2KOH. 

Potassium  hydrosulphide  and  sulphide  precipitate  insoluble  sulphides 
from  the  solutions  of  many  metallic  salts.  They  are  decomposed  by  acids 
with  liberation  of  hydrogen  sulphide. 

When  the  aqueous  solution  of  the  sulphide  is  boiled  with  sulphur  the 
poly  sulphides,  K2S3,  K^S^,  and  K2S-,  are  formed.  The  aqueous  solutions 
of  the  polysulphides  are  decomposed  by  acids,  with  disengagement  of 
hydrogen  sulphide  and  separation  of  sulphur  (milk  of  sulphur).  The  so- 
called  liver  of  sulphur  (^Hepar  sulphiiris,  K.  sulphuratumf  a liver-brown 
mass,  used  in  medicine,  is  obtained  by  the  fusion  of  potassium  carbonate 
with  sulphur,  and  consists  of  a mixture  of  potassium  trisulphide  with 
potassium  sulphate  and  hyposulphite. 

The  aqueous  .solution  of  the  potassium,  as  well  as  that  of  the  sodium 
sulphide,  dissolves  some  metallic  sulphides,  and  forms  sulpho-salts  with 
them  (pp.  223,  225). 


When  dry  ammonia  is  conducted  over  heated  potassium,  potassamide 
(NHjK)  results.  This  is  a dark-blue  liquid  which  solidifies  to  a white, 
crystalline  mass.  It  sublimes  at  about  400°,  and  above  that  temperature 
breaks  down  into  its  elements.  Water  decomposes  it  into  potassium 
hydroxide  and  ammonia. 


Recognition  of  the  Potassium  Compounds. — In  all  of  its  com- 
pounds potassium  is  present  as  a positive  univalent  element.  Almost  all  of 
the  potassium  compounds  are  easily  soluble  in  water.  The  few  exceptions 
serve  for  the  characterization  and  separation  of  potassium.  Tartaric  acid 
added  to  the  solution  of  a potassium  salt  gives  a crystalline  precipitate 
of  acid  potassium  tartrate.  Platinic  chloride  (PtClJ  produces  in  potas- 
sium solutions  a yellow,  crystalline  precipitate  of  PtCl^.  2KCI  (p.  271). 
Potassium  silicofluoride,K2vSiFlp,  is  also  sparingly  soluble  and  can  be  used 
in  detecting  and  estimating  potassium.  Potassium  compounds  introduced 
into  the  flame  of  an  alcohol  or  a gas  lamp  impart  to  the  same  a violet  col- 
oration. The  spectrum  of  the  flame  is  characterized  by  two  bright  lines, 
one  red  and  one  violet  (see  Spectrum  Analysis). 


284 


INORGANIC  CHEMISTRY. 


2.  RUBIDIUM  AND  Ci^:SIUM. 

Rl)  85.4.  Cs  133. 

Rubidium  and  caesium  are  tlie  ])erfcct  analo^rues  of  potassium  (p.  273).  d'hey  were 
discovered  by  means  of  the  spectroscope,  by  lUmsen  and  Kircldioff,  in  i860.  Although 
only  occurring  in  small  quantities,  they  are  yet  very  widely  distributed,  and  fre(juently 
accompany  potassium  in  mineral  sj^rings,  salt,  and  plant  ashes.  'I'he  mineral  lepidolite 
contains  0.5  percent,  of  rubidium;  upward  of  30  i)er  cent,  of  caesium  oxide  is  present 
in  the  very  rare  pollucite,  a silicate  of  aluminium  and  cnesium.  Stassfurt  carnallite  also 
contains  rubiilium.  The  spectrum  of  rubidium  is  marked  by  two  red  and  two  violet 
lines;  caesium  by  two  distinct  blue  lines;  hence  the  names  of  the.se  elements  {riibiilns, 
dark  red  ; itcsiiis,  sky-blue). 

Rubidium  and  caesium  form  double  chlorides  (PtCl^.  2RbCl,  PtCl^.CsCl)  with  plati- 
num chloride,  and  they  are  more  insoluble  than  the  double  platinum  salt  of  potassium, 
hence  may  answer  for  the  separation  of  these  elements  from  potassium,  d'his  is  also 
true  of  the  comi)ounds  2RbCl.SnCl^,  2CsCl.SnCl^  and  2SbCl3.  3CSCI,  which  are  rather 
sparingly  .soluble. 

Rubidium  and  caesium  may  be  obtained  free  by  decomposing  the  fused  chloride  with  the 
electric  current.  Erdmann  and  Kothner  [Ann.  Ch.  294  (1897),  58]  obtained  large  yields  of 
rubidium  by  heating  its  hydroxide  with  magnesium  in  a seamless,  knee-shaped  iron  tube  : 

2RbOII  -(-  2jMg  = 2Rb  2MgO  II.^. 

Dry  hydrogen  is  conducted  through  the  tube  while  heating.  The  rubidium,  which  dis- 
tils over,  is  collected  under  liquid  paraffin.  Metallic  rubidium  has  a silver-white  color, 
with  a .somewhat  yellowish  tinge  ; its  vapor  is  greeni.sh-blue.  Its  .specific  gravity  equals 
1.52;  its  melting  point  is  38.5°.  Oxygen  at  the  ordinary  temperature  converts  it  into 
rubidium  dioxide,  Rb02,  consisting  of  dark-brown  crystals,  which  are  transposed  by 
water  at  a gentle  heat  into  oxygen  and  rubidium  hydroxide  : 

2Rb02  + 2H2  = 2RbOII  -f  H2O  + O. 

The  dioxide  dissolves  in  water  with  hissing  and  the  tumultuous  evolution  of  oxygen  ; 
hydrogen  peroxide  is  produced  at  the  same  time.  The  oxide  RbO  is  not  known.  The 
iodide  at  present  is  quite  frequently  substituted  in  medicine  for  potassium  iodide. 

Ccesium  like  rubidium  is  isolated  by  heating  coesium  hydroxide  with  magnesium  powder 
in  an  atmosphere  of  hydrogen  or  by  the  electrolysis  of  a mixture  of  caesium  and  barium 
cyanides.  Electrodes  of  aluminium  are  employed  for  this  purpose.  Caesium  is  a silver- 
white  metal,  of  specific  gravity  1.85.  It  oxidizes  quite  readily  and  inflames  in  the  air. 
It  melts  at  26.5°  and  boils  at  270°. 

Since  1892  Wells  and  Wheeler,  and  also  Erdmann,  have  prepared  an  interesting 
series  of  caesium  and  rubidium  halides.  The  metals  in  these  appear  to  be  trivalent  and 
also  quinquivalent,  e.  g.,  RbClBr.^,  RbBi^,  RbCl^I,  CsBrg,  CsBr3l2,  Cslj.  Rubidium 
iodine  tetrachloride,  RblCl^,  results  in  conducting  chlorine  into  a rubidium  iodide  solu- 
tion. It  crystallizes  in  monoclinic  yellow  leaflets.  It  dissolves  with  difficulty  in  water. 
Its  solution  acts  as  a powerful  oxidant ; it  dissolves  gold  and  platinum.  It  is  still  undeter- 
mined whether  these  are  atomic  or  only  double  compounds  {e.g.,  RbCl.IClg).  The 
rubidium  and  caesium  compounds  are  distingui.shed  from  those  of  potassium  by  entering 
into  union  more  readily  and  in  greater  proportion  with  other  halides  to  form  double  salts, 
e.  g.,  with  A5CI3,  AsBr3  and  ASI3,  with  which  potassium  double  halides  have  not  been 
prepared. 


3.  SODIUM. 

Na  = 23.05. 

Sodium  is  widely  distributed  in  nature,  especially  as  chloride  in  sea- 
water and  as  rock-salt;  and  isalso  found  in  silicates.  Its  nitrate  is  Cliili 
saltjteter,  and  its  fluoride  in  union  with  aluminium  fluoride  constitutes 


SODIUM. 


285 


the  cryolite  of  Greenland.  The  metal  was  obtained  in  1807,  by  Davy, 
by  the  action  of  a strong  electric  current  upon  fused  sodium  hydroxide. 
It  was  formerly  (1855),  like  potassium,  obtained  upon  a large  scale  by 
igniting  a mixture  of  sodium  carbonate,  finely  divided  anthracite  and 
limestone  in  an  iron  retort : 

Na^COj  -{-20  = 2Na  -[-  3CO. 

A great  advance  was  made  in  its  mansfacture  when  Castner  (1866),  at  Oldbury,  Bir- 
mingham, reduced  the  hydrate  instead  of  the  carbonate  with  carbon  impregnated  with 
finely  divided,  spongy  iron.  Gay-Lussac  and  Thenard  (1808)  had  reduced  the  hydroxide 
with  metallic  iron  at  a white  heat  : 

2NaOH  -j-  2Fe  = 2Na  -j-  2FeO  -)-  H2. 

Netto  (1898,  Wallsend,  Newcastle-on-Tyne)  invented  a most  satisfactory  process, 
abandoning  it,  however,  later.  It  consisted  in  allowing  molten  caustic  soda  to  run  over 
wood  charcoal  placed  in  vertical  iron  retorts.  Sodium  vapors  escaped  constantly  from 
an  exit  tube  at  the  top,  while  fused  carbonate  ran  out  from  a tube  at  the  bottom  : 

3NaOH  + C Na2C03  + Na 

The  sodium  vapors  were  condensed  in  flat  iron  receivers  and  the  liquid  metal  was 
collected  under  rock  oil. 

Castner  and  Kellner  electrolyzed  a salt  solution,  using  mercury  as  anode  ; the  amalgam 
was  then  distilled  when  mercury  was  expelled  and  sodium  remained. 

Grabau  electrolyzed  fused  .salt  after  reducing  its  melting  point  by  adding  potassium 
and  strontium  chlorides.  The  sodium  then  escaped  in  vapor  form. 

Sodium  can  be  made  on  a small  scale  by  heating  the  peroxide  with  freshly  ignited 
wood  charcoal : 

3Na20.2  -f  2C  = Na.2  + 2Na2C03. 

Calcium  carbide  may  be  substituted  for  the  charcoal  : 

7Na202  + 2CaC2  = 2CaO  -f  4Na2C03  -|-  3Na2. 

The  action  in  both  cases  is  very  energetic. 

Sodium  in  external  properties  is  very  similar  to  potassium.  It  melts  at 
95.6°,  boils  at  742°,  and  is  converted  into  a colorless  vapor,  which 
burns  with  a bright  yellow  flame  in  the  air.  It  oxidizes  readily  on  ex- 
posure, and  decomposes  water  even  in  the  cold,  although  less  energet- 
ically than  potassium.  A piece  of  sodium  thrown  upon  water  swims 
about  upon  the  surface  with  a rotary  movement,  the  disengaged  hydrogen, 
however,  not  igniting.  If  we  prevent  the  motion,  by  confining  the 
metal  to  one  place,  the  heat  liberated  by  the  reaction  attains  the  ignition 
temperature  of  hydrogen,  and  a flame  follows,  as  was  also  observed  with 
p jtassium  (p.  40). 

Sodium  Oxide,  NajO,  is  not  definitely  known. 

Sodium  Peroxide,  Na202,  has  recently  been  introduced  into  commerce  as  a 
bleaching  agent.  It  is  made  by  heating  sodium  in  a stream  of  dry  air,  using  vessels  of 
aluminium  and  a temj)erature  below  300°.  It  is  a yellow-white  ]X)wder.  It  melts  with 
greater  difficulty  than  caustic  .soda,  and  at  elevated  temperatures  gives  off  oxygen.  Water 
decomposes  it  with  caustic  soda  and  oxygen  which  escapes  uj)()n  boiling.  With  ice-water 
it  forms  a solution  containing  both  caustic  .soda  and  hydrogen  j)eroxide.  On  careful  evapora- 
tion such  solutions  yield  crystalline  hydrates  of  .sodium  superoxide.  Anhydrous  acids  and 
alcohol  appear  to  decompose  it  with  the  formation  of  a peculiar  hydrate  : 

NajO,  i IlCl  = NaCl  + Na02H, 


286 


INORGANIC  CHEMISTRY. 


wliich  probably  lias  the  formula  Na-()-0-H.  Sodium  superoxide  acts  on  many  organic 
compounds  with  the  production  of  flame  and  the  separation  of  carbon.  At  a red  heat  it 
is  superior  to  all  other  oxidants  in  its  powerful  action  (i  jiart  of  the  material,  2 parts  of 
soda  and  4 parts  of  .sodium  superoxide). 

Sodium  Hydroxide,  sodium  hydrate,  or  caustic  soda,  NaOH,  like 
])Otassium  hydroxide,  is  formed  by  boiling  a solution  of  soditim  car- 
bonate with  calcium  hydroxide: 

Na.^COg  + Ca(OII)2  = CaCOj  + 2NaOII. 

At  present  it  is  directly  protluced  in  the  soda  manufacture  by  adding  a 
little  more  carbon  to  the  fusion  (see  Soda);  or  by  the  electrolysis  of 
sodium  chloride  (jtp.  290,  291). 

'bhe  sodium  hydroxide  which  solidifies  after  fusion  is  a white,  radiating, 
crystalline  mass,  and  re.sembles  caustic  ]iotash  very  much.  It  attracts 
water  from  the  air,  becomes  moist,  and  coats  itself  by  carbon  dioxide 
absorption  with  a white  layer  of  sodium  carbonate  (caustic  potash  del- 
i(piesces  perfectly,  because  the  resulting  carbonate  is  al.so  deliquescent), 
d'he  aqueous  solution,  called  sodium  hydroxide,  resembles  that  of  potas- 
sium. Crystals  of  NaOH  -j-  separate  at  0°  from  the  concentrated 

solution  ; they  melt  at  6°. 

Sodium  Chloride,  NaCl,  is  abundant  in  nature.  It  is  found  almost 
everywhere  in  the  earth  and  in  natural  waters;  in  sea-water  it  averages 
2.  7-3. 2 per  cent.  As  rock-salt  it  forms  large  deposits  in  many  districts, 
especially  at  Stassfurt  and  Wieliczka  in  Galicia  (p.  276). 

In  warm  climates,  on  the  coasts  of  the  Mediterranean  Sea,  sodium  chloride  is  gotten 
from  the  sea  according  to  the  following  procedure  : At  high  tide  sea- water  is  allowed 
to  flow  into  wide,  flat  basins  (salt  gardens),  in  which  it  evaporates  under  the  sun’s  heat ; 
the  working  is  limited,  therefore,  to  summer  time.  After  sufficient  concentration,  pure 
sodium  chloride  first  separates,  and  this  is  collected  by  itself.  Later,  there  crystallizes  a 
mixture  of  sodium  chloride  and  magnesium  sulphate  ; finally  potassium  chloride,  magne- 
sium chloride  and  some  other  salts  appear  (among  them  potassium  iodide  and  bromide), 
the  .separation  of  which  constitutes  a special  industrial  branch  in  some  regions.  In 
cold  climates,  as  in  Norway  and  at  the  White  Sea,  the  cold  of  winter  is  employed  for  the 
production  of  salt.  In  the  freezing  of  sea-water,  as  well  as  of  other  solutions,  almost 
pure  ice  separates  at  first ; the  enriched  sodium  chloride  solution  is  then  concentrated  in 
the  usual  way. 

Rock-salt  is  either  mined  in  shafts,  or,  where  the  strata  are  not  so  large  and  are 
admixed  with  other  varieties  of  rock,  alixiviation  process  is  employed.  Borings  are  made 
in  the  earth  and  wmter  runs  into  them,  or  into  any  openings  already  formed.  When  the 
water  has  saturated  itself  with  sodium  chloride,  it  is  pumped  to  the  surface  and  the  brine 
then  further  worked.  In  many  regions,  especially  in  Reichenhall,  in  Bavaria,  more  or 
less  saturated  natural  salt  or  brine  springs  flow  from  the  earth.  The  concentration  of  the 
non-saturated  brine  occurs  at  first  in  the  so-called  “graduation”  houses.  These  are  long 
wooden  frames  filled  with  fagots,  and  on  letting  the  salt  water  run  upon  them  it  will  be 
distributed  and  evaporated  by  the  fall  ; the  concentrated  brine  collects  in  the  basin  below, 
and  is  then  evaporated  over  a free  fire. 

Sodium  chloride  crystallizes  from  water  in  transparent  cubes,  of  specific 
gravity  2.13,  which  arrange  themselves  by  slow  cooling  into  hollow,  four- 
sided pyramids.  It  melts  at  8 15°  and  volatilizes  at  a white  heat.  It  is  not 
much  more  soluble  in  hot  than  in  cold  water  ; 100  parts  at  0°  dissolve  36 
l)arts  of  salt ; at  100”,  39  ])arts.  The  saturated  solution,  therefore,  contains 


SODIUM  SULPHATE. 


287 


about  26  per  cent,  of  sodium  chloride.  If  the  saturated  solution  be  cooled 
below  — 10,  large  monoclinic  tables  (NaCl  -f-  2 H,^0)  separate  ; these  lose 
water  at  0°  and  become  cubes. 

The  ordinary  sodium  chloride  usually  contains  a slight  admixture  of 
magnesium  salts,  in  consequence  of  which  it  gradually  deliquesces  in  the 
air  ; the  perfectly  pure  salt  is  not  hygroscopic.  When  heated  the  crystals 
crackle,  because  of  the  escape  of  the  mechanically  enclosed  water. 


Sodium  Bromide  and  Iodide  crystallize  at  ordinary  temperatures  with  two  molecules 
of  water,  which  they  lose  again  at  30°  ; above  30°  they  separate  in  anhydrous  cubes. 
Sodium  bromide  fuses  at  760°  and  the  iodide  at  690°  ; the  former  is  difficultly  soluble  in 
alcohol  and  the  latter  is  very  soluble. 

Sodium  Chlorate  (NaC103)  and  Perchlorate  (NaClO^)  are  considerably  more  solu- 
ble in  water  than  the  corresponding  potassium  salts. 

Sodium  lodate,  NalCXj,  is  obtained  in  the  same  manner  as  the  potassium  salt,  and 
at  ordinary  temperatures  crystallizes  with  three  molecules  of  water  in  silky  needles.  It 
is  pre.sent  in  Chili  saltpeter.  If  chlorine  gas  be  conducted  through  the  warmed  solution 

of  sodium  iodate  in  sodium  hydroxide,  the  periodate  10  <!  P-  crystal- 

lizes  out  on  cooling.  This  becomes  the  normal  salt  (NalO^  -|-  3H2O)  when  dissolved  in 
nitric  acid. 


Sodium  Sulphate  (A^atrmm  sulphuriciiiJi),  Na2SO^,  crystallizes  at 
ordinary  temperatures  with  ten  molecules  of  water  of  crystallization,  and 
is  then  known  as  Glauber’s  ’idXx.  {Sal  mirabile  Glauberi').  It  occurs  in 
many  mineral  waters,  and  in  large  deposits,  with  or  without  water  of 
crystallization,  in  Spain.  It  is  a by-product  in  the  manufacture  of 
sodium  chloride  from  sea-water  and  brine.  It  is  produced  in  large  quan- 
tities by  heating  salt  with  sulphuric  acid  : 

2NaCl  + H2SO,  = Na^SO,  -f  2HCI, 

and  is  used  in  making  soda  (sodium  carbonate).  Or  it  may  be  prepared 
by  the  method  of  Hargreaves,  by  conducting  sulphur  dioxide,  air  and 
steam  over  strongly  ignited  sodium  chloride: 

2NaCl  + SO,  -f  O T H,0  Na-^SO,  + 2HCI. 

More  recently  the  sulphate  has  been  obtained  by  a transposition  of 
sodium  chloride  with  magnesium  sulphate  at  a winter  temperature — a 
procedure  which  is  prosecuted  chiefly  in  Stassfurt,  where  immense  quan- 
tities of  magnesium  sulphate  (kieserite)  exist : 

2NaCl  + MgSO,  = MgCl,  + NaaSO,. 

Sodium  sulphate  crystallizes  at  ordinary  temperatures  with  ten  molecules 
of  water,  in  large,  colorless,  monoclinic  prisms,  which  crumble  in  the  air 
and  fall  into  a white  powder.  The  salt  Na2SO^  -j-  10H2O  is  only  stable  at 
temperatures  below  34°.  It  melts  at  34°,  separating  into  a saturated 
solution  and  the  anhydrous  salt  Na2SO^.  The  latter  melts  at  about  886°. 

This  explains  the  remarkable  alteration  which  the  .solubility  of  Glauber’s  salt  sustains 
in  water  at  34°.  As  a rule  the  solubility  of  a substance  in  water  does  not  change  sud- 


288 


TNOUGANIC  CHRMISTRY. 


flenly  with  the  temperature,  l)iit  f,Mnflually.  A siulch'ii  cliaufre  in  .sohihilify  is  conuccled 
witli  a chanj^e  in  coTidition  of  the  substance  passinjr  into  soiutif)n.  'I’liis  apj)lies  also  to 
(dauber’s  salt.  lOO  jiarts  of  water  dissolve,  at  o°,  12  parts;  at  18°,  48  parts  ; at  25°, 
too  parts;  at  30°,  200  parts;  at  34°,  354  ])arts  of  the  hydrous  .salt.  At  the  last  tem- 
perature the  solubility  is  f^reatest  ; by  further  increase  of  heat  it  ^rradually  diminishes;  at 
50°,  ICX)  parts  of  water  (lissolve  only  263  parts  ; at  lOO°,  238  parts  of  the  salt.  'I'his  is 
explained  l)y  the  fact  tliat  up  to  34°  the  solubility  relates  to  the  salt  Na.^SO,  ioII.^(). 
whicli  cannot  exist  at  higher  temperatures.  From  this  point  upwards  the  solubility 
relates  to  the  anhydrous  salt  wliich  separates  as  a deposit  in  small  rhombic  pyramids,  if 
the  solution  of  the  hydrous  salt,  saturated  at  34°,  be  healed  higher. 

When  the  .solution,  saturated  at  34°,  is  allowed  to  cool  to  the  ordinary  tem])erature, 
and  even  lower,  not  the  slightest  sei)aration  of  crystals  occurs,  although  the  salt  is  vastly 
more  in.soluble  at  lower  temperatures  than  at  34°.  Many  other  .salts  form  snpersaluraled 
solutions,  although  they  are  less  striking  than  that  of  (dauber’s  salt,  d'he  supersaturated 
solution  of  the  latter  may  be  agitated  and  twirled  about  without  cry.stallization  setting  in. 
If,  however,  a glass  rod,  or  some  other  .solid  body,  be  introduced  into  the  .solution,  it  will 
solidify  suddenly  to  a crystalline  mass,  d'he  particles  of  dust  floating  about  in  the  air 
will  have  a like  effect ; therefore,  to  preserve  the  supersaturated  solution,  the  ve.ssel  con- 
taining it  should  be  kept  well  corked.  Ily  accurately  made  investigations,  it  has  been 
determined  that  the  crystallization  of  the  supersaturated  (dauber’s  .salt  solution  is  only 
induced  by  contact  with  already  formed  crystals.  These  must  then  be  pre.sent  every- 
where in  the  atmosphere,  because  only  solids  that  have  been  expo.sed  to  the  air,  and 
that  have  not  been  carefully  cleansed  afterward,  bring  about  the  crystallization. 

In  the  crystallization  of  a supersaturated  (dauber’s  salt  .solution  considerable  heat  is 
disengaged,  and  the  mass  increases  in  tem{)erature.  This  is  becau.se  the  latent  heat  of  all 
substances  in  the  liquid  condition  is  greater  than  in  the  solid.  At  io°,  occasionally,  and 
of  their  own  accord,  tran.sparent  crystals,  Na2SO^  71^2^^  .separate  from  the  supersatu- 
rated solutions.  These  crystals  change  readily  to  anhydrous  .sodium  sulphate  and 
Glauber’s  salt. 

This  salt  is  employed  in  medicine  as  a purgative,  and  finds  extended 
application  in  the  manufacture  of  glass  and  in  the  preparation  of  soda. 

The  primary  or  acid  sodium  sulphate,  NaHSO^,  is  obtained  by 
the  action  of  sulphuric  acid  upon  the  neutral  salt  or  upon  sodium  chloride  : 


NaCl  + H2SO,  = NaHSO,  -f  fICl. 


At  ordinary  temperatures  it  crystallizes  with  one  molecule  of  water, 
and  is  perfectly  analogous  to  the  potassium  salt. 

The  sodium  salts  of  sulphurous  acid  are  obtained  by  conducting  sulphur  dioxide  into 
solutions  of  sodium  hydroxide  or  carbonate.  The  secondary  sulphite,  Na2S03,  crystallizes 
with  seven  molecules  of  water  at  ordinary  temperatures  ; in  the  presence  of  sodium 
hydroxide,  or  by  warming  the  solution,  it  separates  in  the  anhydrous  state.  The  prunary 
sulphite,  NallSOa,  consisting  of  small,  easily  .soluble  crystaks,  gives  up  sulphur  dioxide  in 
the  air,  and  is  oxidized  to  sodium  .sulphate.  According  to  Schwicker,  when  it  is  neutralized 
with  potassium  carbonate  it  forms  yellow,  indistinct  crystals  of  sodium  potassium  sul- 
phite, NaKSO.j  -(-  II-iG.  Primary  potassium  sulphite  on  the  other  hand,  when  neutral- 
ized with  .soda,  yields  the  salt  NaKSO^ -p  2H2O,  consisting  of  hard,  yellowish  crystals. 
'I'his  forms  entirely  different  compounds  with  organic  iodides  from  those  which  the  first 
.salt  (with  which  it  is  i.someric,  if  the  water  of  cry.stallization  be  di.sregarded)  affords.  Two 
isomeric  .sodium  potassium  hyposulphites,  NaKS2()3,  are  obtained  when  these  salts  take 
up  sulphur.  'I'he  following  formulas  express  their  structure  : 


K_S(  )2-(  )Na 

KS-SeVONa 


Na-SC)2-OK 

NaS-S02-OK. 


[Her.  22  (1889),  1728;  .see  also  pp.  186,  198.] 


SODIUIM  CARBONATE. 


289 


Sodium  Hyposulphite,  Na2S203,  Sodium  Thiosulphate,  is  pre- 
pared by  boiling  the  aqueous  solution  of  neutral  sodium  sulphite  with 
flowers  of  sulphur  : 

Na2S03  -j-  S = Na2S203. 

It  is  obtained  as  a by-product  in  the  recovery  of  sulphur  from  the  soda 
residues.  It  crystallizes  with  five  molecules  of  water,  in  large  monoclinic 
prisms,  dissolves  very  readily  in  water,  and  is  somewhat  deliquescent  in 
the  air.  At  56°  it  melts  in  its  water  of  crystallization  ; loses  all  water 
at  100°,  and  decomposes  by  further  heating  into  sodium  sulphate  and 
sodium  pentasulphide,  Na2S..  When  the  dry  salt  is  heated  in  the  air,  the 
sulphur  of  the  polysulphide  burns  with  a blue  flame.  Acids  decompose 
the  aqueous  solution  with  separation  of  sulphur  and  evolution  of  sulphur 
dioxide  (p.  197)  : 

Na2S203  -f  2HCI  = 2NaCl  -f  SO2  + S -f  H2O. 

Like  the  sulphate,  it  readily  affords  supersaturated  solutions.  The 
hyposulphite  is  used  as  a reducing  agent ; chlorine,  bromine  and  iodine 
are  converted  by  it  into  the  corresponding  halogen  salts.  An  iodine 
solution  is  instantaneously  decolorized  by  sodium  hyposulphite  with  the 
production  of  sodium  tetrathionate  (p.  199).  Chlorine  behaves  differ- 
ently; sulphuric  acid  and  sodium  chloride  are  produced.  Upon  this 
reaction  rests  the  application  of  sodium  hyposulphite  as  an  mitichlor  in 
chlorine  bleaching,  to  remove  the  excess  of  the  chlorine,  which  has  a 
destructive  action  upon  the  fiber.  In  consequence  of  its  property  of  dis- 
solving the  halogen  silver  derivatives,  it  is  employed  in  photography. 

Sodium  Carbonate  (Soda),  Na2C03.  This,  technically,  very  im- 
portant salt  is  obtained  (i)  as  natural  soda,  (2)  from  the  ashes  of  plants, 
and  (3)  by  chemical  methods. 

1.  Natural  soda  disintegrates  from  the  soil  in  certain  districts,  as  in 
Hungary,  Asia,  and  Africa,  during  the  hot  seasons.  It  is  a constituent 
of  many  mineral  springs,  e.  g.,  Carlsbad,  and  of  the  soda  seas  of  Egypt, 
of  Central  Africa,  and  of  the  coasts  of  the  Caspian  and  Black  Seas,  of 
California,  and  of  German  East  Africa.  In  these  seas  or  lakes  a salt 
deposits  having  the  formula 

Na2C03  -h  NallCO,  + 2H2O. 

In  Egy])t  it  is  called  iro-tia  and  in  Colombia  urao.  Its  occurrence  at 
Owen’s  Lake,  Inyo  County,  California,  is  the  only  one  of  scientific 
importance.  The  waters  of  this  lake  upon  evaporation  yield  very  large 
quantities  of  quite  pure  soda. 

2.  Soda  from  the  Ashes  of  Plants. — Before  Le  Blanc’s  process  came 
into  use  the  greater  ])ortion  of  soda  was  made  from  the  ashes  of  sea  and 
coast  \)\diY\X.?>  (^Chefiopodium,  Salsola,  Atriplex,  Salicornia,  etc.);  these 
assimilate  the  sodium  salts  of  the  earth,  while  the  land  plants  absorb  the 
potassium  salts,  and  for  this  reason  contain  potashes  in  their  ash.  In 
Southern  France  and  in  Spain  beach  ])lants  were  and  are  yet  specially 
cultivated  for  this  purpose.  The  Spanish  soda  was  particularly  rich  in 

25 


290 


INORGANIC  CHEMISTRY. 


carbonate  and  controlled  the  markets  for  a long  period.  At  present  soda 
from  the  ashes  of  plants  possesses  only  a local  interest, 

3.  In  the  chemical  way  soda  is  pre[)ared  almost  exclusively  from  salt 
(NaCl).  Tlie  following  methods  are  in  use:  (i)  That  discovered  in 
1 794  by  Le  Blanc;  (2)  the  ammonia-soda  i)rocess  introduced  in  1866,  and 

(3)  the  electrolytic  method  of  recent  date.  To  these  may  be  added  the 

(4)  cryolite  soda  process, 

I.  In  the  Le  Blanc  method  the  sodium  chloride  is  converted  into 
sodium  sulphate  by  warming  with  sulphuric  acid  (pj).  58,  287).  When 
the  latter  is  dry,  it  is  mixed  with  charcoal  and  chalk  (calcium  carbonate) 
and  ignited  in  a reverberatory  furnace,  d'wo  principal  ])hases  may  be 
distinguished  in  this  reaction.  First,  the  carbon  reduces  the  sodium 
sulphate  to  sodium  sulphide: 

Na,SO,  + 2C  Na,S  f-  2CO,. 

The  sodium  sulphide  then  acts  upon  the  calcium  carbonate  to  form 
calcium  sulphide  and  sodium  carbonate: 

Na.^S  -f-  CaCOj  = CaS  -(-  Na^COg. 

At  the  same  time  the  high  temperature  converts  a portion  of  the  cal- 
cium carbonate  into  calcium  oxide  and  carbon  dioxide,  which  is  reduced 
by  the  ignited  carbon  to  the  monoxide: 

CaC03  = CaO  + CO, ; CO,  + C = 2CO. 

The  appearance  of  the  latter,  which  burns  with  a bluish  flame,  indicates 
the  end  of  the  action.  The  chief  products  in  the  soda  fusion  are,  then, 
sodium  carbonate  and  calcium  sulphide,  mixed  with  varying  amounts  of 
calcium  carbonate,  calcium  oxide,  and  foreign  substances.  This  fused 
mass  is  called  C7'iide  soda.  It  is  lixiviated  in  specially  constructed  appa- 
ratus with  cold  water;  the  sodium  carbonate  dissolves,  and  there  remain 
behind  calcium  sulphide,  calcium  carbonate,  and  a portion  of  the  foreign 
substance — the  soda  residue. 

During  the  lixiviation  the  caustic  lime  present  in  the  crude  soda  acts 
upon  the  sodium  carbonate  with  the  assistance  of  water,  and  there  result 
calcium  carbonate  and  sodium  hydrate.  The  latter  passes  into  solution 
with  the  soda : 

CaO  T + Na2C03  = CaCOg  + 2NaOH. 

It  is  possible  by  the  Le  Blanc  process  to  obtain  a preponderance  of 
scKlium  hydroxide  if  the  soda  fusion  be  mixed  at  the  beginning  with  more 
carbon,  heated  intensely  and  the  crude  soda  be  then  extracted  with  hot 
water. 

When  the  solution  is  evaporated  the  soda  will  separate  from  the  hot 
lirpiid  as  a crystalline  ])owder,  Na^COg  -j-  H.^O.  It  is  removed  from  the 
li(|uid,  and  new  licpiors  are  introduced,  etc.  The  mother  liquor,  the 
so-called  red  liipior,  contains  finally  caustic  soda  and  sodium  suli)hide 
almost  exclusively.  The  soda  flour  is  freed  from  the  mother  licpior  by  a 
centrifugal,  dried  and  calcined-— soda.  For  further  purification 


SODIUM  CARBONATE. 


291 


it  is  recrystallized  from  water,  when  it  separates  in  large,  transparent 
crystals  of  the  formula  Na.^COg  -j-  10H2O — crystallized  soda. 

The  by-products  and  the  refuse  in  the  manufacture  of  soda — hydrochloric  acid  and 
soda  residues — must  be  utilized  as  fully  as  possible,  because  of  the  enormous  competition 
encountered  by  the  manufacturers.  To  this  end  the  hydrochloric  acid  is  converted  by  the 
process  described  on  page  50,  into  chlorine  and  bleaching  lime  and  the  soda  residues  are 
worked  in  various  ways  to  get  their  sulphur  content  into  an  available  form.  Of  late 
years  the  Chance-Claus  method  has  been  adopted.  It  consists  in  decomposing  the 
residues  with  carbonic  acid:  CaS -(-  H.^O -f- CO.^  = CaCO.^  I l.,S,  and  burning  the 
liberated  hydrogen  sulphide  either  with  an  insufficiency  of  air,  when  sulphur  will  sepa- 
rate : H.^S  -|-  O = H.^O  + S,  or  with  an  excess  of  air  when  sulphur  dioxide  is  formed  : 
H^S  -r  3O  = 11.2^^  + SO.^.  The  last  product  is  then  conducted  into  lead  chambers 
(p.  189).  About  70,000  tons  of  sulphur  are  recovered  annually  (regenerated  sulphur). 

2.  The  ammonia-soda  process  is  based  upon  the  transposition  of 
sodium  chloride  with  })rimary  ammonium  carbonate  to  ammonium 
chloride  and  primary  sodium  carbonate  : 

NaCl  ^ NHdICOg  ^NallCOg  T NH^Cl. 

This  change  takes  place  at  the  ordinary  temperature.  The  acid  sodium 
carbonate  being  sparingly  soluble  in  cold  water,  is  converted  into  sodium 
carbonate  upon  ignition  : 

2NaHC03  ^ Na^COg  + CO,  + H,0. 

The  ammonium  chloride  remains  in  solution.  Ammonia  is  recovered 
from  it  by  means  of  lime.  In  actual  practice  carbon  dioxide  is  con- 
ducted under  pressure  into  a concentrated  salt  solution  saturated  with 
ammonia : 

NaCl  -f  NHg  + CO,  -f  H^O  = NH^Cl  -f  NaHCO^. 

The  temperature  must  not  exceed  40°.  The  carbonic  acid  is  obtained 
by  burning  lime,  and  when  the  bicarbonate  is  heated  half  of  it  is 
recovered.  The  lime  from  the  calcite  is  used  to  generate  ammonia  from 
ammonium  chloride.  The  raw  material  in  this  process  consists  therefore 
of  salt  and  limestone  which,  with  the  assistance  of  ammonium  salts,  are 
converted  into  soda  and  calcium  chloride  (p.  50).* 

This  process  is  exceedingly  simple  from  the  chemical  standpoint. 
Difficulties  arose  in  building  the  neces.sary  ajiparatus  on  a technical  scale 
and  they  militated  against  the  general  adoption  of  the  method.  Fortu- 
nately they  have  lieen  completely  overcome,  and  this  is  due  in  a large 
measure  to  E.  Solvay,  so  that  at  present  the  ])roduction  of  soda  by  the 
Le  Blanc  yirocess  is  becoming  less  frequent.  It  is  only  in  England  that 
the  old  method  holds  sway. 

3.  The  Le  Blanc  method  with  it.s  improvement.s  was  in  operation  for  a century  before 
it  was  displaced  by  the  ammonia-soda  process,  and  now  the  latter  is  being  .seriously 
threatened  by  a dangerous  rival — the  electrolytic  production  of  caustic  alkalies,  alkaline  car- 
bonates, chlorine  and  potassium  chlorate  ( p.  275).  When  an  af|ueous  salt  solution  is  elec- 
trolyzed chlorine  .separates  at  the  anode  and  sodium  at  the  kathode.  'I'he  latter  acts 


* Magnesite  and  magnesia  can  be  substituted  for  limestone,  and  lime. 


292 


INORGANIC  CMKMISTRY. 


immediately  on  tlie  water,  producing  hydrogen  and  sodium  hydroxide,  d'lie  free  chlorine 
would  produce  sodium  chloride  and  hypochlorite  or  chlorate  (pp.  175,  176)  if  the  katluxle 
and  anode  licpiors  were  not  separated  by  a porous  diaj)hragm.  1 tifficulties,  in  the  construct- 
ing of  the  diaidiragms,  confronted  the  technical  utilization  of  the  process.  'There  is  nosul)- 
stance  which  does  not  gradually  disintegrate  when  used  as  anode.  However,  all  these 
objectionable  features  have  been,  in  a measure,  overcome.  'The  sodium  hydroxide  is 
obtained  as  such  or  it  is  separated  in  the  form  of  the  sparingly  soluble  bicarbonate  ui)on 
conducting  carbonic  acid  through  its  solution.  'The  chlorine  is  brought  into  trade  in  the 
liquid  form  or  it  is  changed  to  bleaching  lime. 

4.  Considerable  cjuantities  of  soda  are  obtained  at  present  from  cryolite,  a compound 
of  aluminium  fluoride  and  sodium  fluoride  (AlT'l.,.  3NaFl),  which  occurs  in  great  deposits 
in  Greenland.  The  pulverized  mineral  is  ignited  with  burned  lime  or  chalk  ; insoluble 
calcium  fluoride  and  a very  soluble  compound  of  aluminium  oxide  with  sodium  oxide, 
called  sodium  aluminate  (see  Aluminium),  are  produced: 

2(A1F13. 3NaFl)  -f  6CaO  =r  bCaFl^  -f  Al.Og.  3Na20. 

The  mass  is  treated  with  water  and  carbon  dioxide,  obtained  by  burning  lime,  con- 
ducted into  the  solution,  which  causes  the  precipitation  of  aluminium  oxide,  and  sodium 
carbonate  dissolves  : 

A\,0, . 3Na,0  + 3H2O  + 3CO,  = Al, (011)6  + 3Na,C03. 

This  method  is  no  longer  in  use  in  Europe.  It  continues  of  value  in  North  America 
and  in  Denmark,  because  the.se  countries  confrol  large  deposits  of  cryolite.  Calcium 
fluoride  is  largely  employed  by  them  in  the  manufacture  of  glass  and  porcelain  ; the 
aluminium  oxide  is  used  for  making  alum,  aluminium  sulphate  or  metal. 

At  ordinary  temperatures  sodium  carbonate  crystallizes  with  ten  mole- 
cules of  water  (Na2C03 -f- 10H2O)  in  large  monoclinic  prisms,  which 
crumble  upon  exposure  and  become  a white  powder.  It  melts  at  50°  in 
its  water  of  crystallization,  and  upon  heating  a pulverulent  hydrate, 
Na2C03  -j-  2H2O,  separates,  which  in  dry  air  has  one  molecule  of  water, 
and  at  100°  loses  all  of  this.  At  30-50°  rhombic  prisms  of  the  compo- 
sition Na2C03  7H2O,  crystallize  from  the  aqueous  solution.  The  anhy- 
drous salt  absorbs  water  from  the  air  but  does  not  deliquesce.  It  melts 
at  850°  and  volatilizes  somewhat  at  a very  high  temperature.  100  parts 
of  water  dissolve  7 parts  at  0°,  and  at  38°,  52  parts  of  the  anhydrous 
salt.  At  more  elevated  temperatures  the  solubility  is  less,  as  in  the  case 
of  the  sulphate.  Sodium  carbonate  has  a strong  alkaline  reaction  ; acids 
liberate  carbon  dioxide  from  it. 

Primary  Sodium  Carbonate,  ordinary  bicarbonate  of  soda  (Nn- 
triimi  bicarbonicuni),  NaHC03,  produced  by  the  action  of  carbon 
dioxide  upon  the  hydrous  secondary  carbonate  : 

Na2C03  + CO3  + II2O  2NaIlC03. 

It  crystallizes  without  water,  in  small  monoclinic  tables;  it  dissolves, 
however,  at  ordinary  temperatures  in  lo-ii  parts  of  water,  and  possesses 
feeble  alkaline  reaction.  By  heating  and  by  boiling  the  solution  it  ])asses 
into  the  secondary  carbonate  with  disengagement  of  carbon  dioxide, 
'i'he  dry  salt  decomposes  rafiidly  even  below  100°.  By  raj)id  evaporation 
small  moiioclinic  prisms  of  the  so-called  sodium  sesquicarbonate, 
Na2(J03  2NallC03  -j-  2II2O,  separate.  The  salt  which  deposits  in  the 


SODIUM  PHOSPHATES. 


293 


sodium  seas  of  Hungary  and  E'^vpt,  has  the  composition  Na.,CO,. -1- 
NaHCO,  + 2Hfi  (p.  289). 

Sodium  Nitrate,  NaNOg,  Chili  saltpeter,  is  found  in  immense 
deposits  in  Peru.  The  saltpeter  earth  contains  of  sodium  nitrate  from  15 
to  65  parts  out  of  100  jiarts,  the  rest  being  sodium  chloride,  a little  potas- 
sium nitrate,  potassium  perchlorate  and  sodium  iodate  (pp.  55,  i 78).  The 
soil  is  extracted  with  boiling  water  ; on  cooling,  crude  saltpeter  sejiarates ; 
it  is  purified  by  recrystallization.  It  crystallizes  in  rhombohedra  very 
similar  to  cubes,  hence  designated  cubic  saltpeter,  to  distinguish  it  from 
the  prismatic"  potassium  saltpeter.  It  fuses  at  about  318°.  In  water 
it  is  somewhat  more  easily  soluble  than  potassium  saltpeter.  In  the  air  it 
attracts  moisture,  hence  it  is  not  adapted  for  the  manufacture  of  gun- 
powder. In  other  respects  it  is  perfectly  similar  to  potassium  nitrate.  It 
is  largely  used  in  the  manufacture  of  nitric  acid,  and  especially  in  pre- 
paring potassium  saltpeter  (p.  279). 

Sodium  Nitrite,  NaNOg,  is  prepared  like  potassium  nitrite  (p.  280), 
by  heating  sodium  nitrate  with  lead,  iron,  or  graphite.  It  crystallizes 
more  readily  than  potassium  nitrite,  and  does  not  deliquesce  in  the  air. 
It  occurs  in  trade  in  small  colorless  crystals,  containing  from  93  to  98  per 
cent,  of  the  pure  salt.  It  is  largely  used  in  the  dye  industry  for  the 
preparation  of  the  azo- compounds. 

Sodium  Phosphates.— The  sodium  salts  of  phosphoric  acid  are  less 
soluble  and  crystallize  better  than  those  of  potassium.  The  trisodiwji 
phosphate,  NagPO^,  is  made  by  saturating  one  molecule  of  phosphoric  acid 
with  three  molecules  of  sodium  hydroxide,  and  crystallizes  in  six-sided 
prisms  with  twelve  molecules  of  water.  It  has  a strong  alkaline  reaction, 
absorbs  carbon  dioxide  from  the  air,  and  is  converted  into  the  secondary 
salt. 

Disodium  phosphate,  Na2HPO^,  is  the  most  stable  of  the  sodium  phos- 
phates, and  hence  is  generally  employed  in  laboratories  (JSfatriuin  phos- 
phoricwti).  It  may  be  obtained  by  saturating  phosphoric  acid  with 
sodium  hydroxide  to  feeble  alkaline  reaction,  or  may  be  prepared  on  a 
large  scale  by  decomposing  bone  ashes  (tri calcium  phosphate)  with  an 
equivalent  amount  of  sulphuric  acid  and  precipitating  the  calcium  as 
dicalcium  phos])hate  with  soda.  It  crystallizes  at  ordinary  temperatures 
v/ith  twelve  molecules  of  water  in  large  monoclinic  prisms  which  effloresce 
rapidly  in  the  air.  It  separates  from  solutions  with  a temperature  above 
30°  in  non-efflorescing  crystals  containing  seven  molecules  of  water.  It 
is  soluble  in  4-5  parts  of  water,  and  shows  a feeble  alkaline  reaction. 
Wlien  heated  the  salt  loses  water,  melts  at  about  300°  and  becomes 
sodium  pyrophosphate,  Na^P20.j,  which  crystallizes  with  ten  molecules 
of  water,  and  upon  boiling  with  nitric  acid  passes  into  primary  sodium 
phosphate. 

The  primary  or  ?nonosodium  phosphate,  NaH2P04,  crystallizes  with  one 
molecule  of  water,  and  exhibits  a faintly  acid  reaction.  At  100°  it  loses 
its  water  of  crystallization,  and  at  200°  becomes  disodium  pyn'ophosphate, 
Na2H2P.^O^,  which  at  240°  forms  sodium  metaphosphate,  NaPOg : 


Na2lbP,0-  = 2NaP()g  4 IlgC). 


294 


INORGANIC  CHEMISTRY. 


\Vc  get  various  modifications  of  the  inelapliospliate,  according  to  the 
conditions  of  fusing  and  cooling;  they  are  probably  polymerides,  cor- 
res})onding  to  the  torinulas  Na.^P./)^^,  NaJ’.,()y,  etc.  Uiion  heating  sodium 
metaphosphate  with  metallic  oxides  the  latter  dissolve,  and  salts  of  ortho- 
phosphoric  acid  are  formed,  c.  .* 

NaP03  -f  CuO  = CuNaPO,. 

In  this  manner,  characteristic  colored  glasses  (phosyihorus  beads)  are 
obtained  with  various  metals.  In  blowi)ii)e  analysis  this  behavior  serves 
for  the  detection  of  the  respective  metals. 

The  salts  of  arsenic  acid  i)erfectly  analogous  to  those  of  phosphoric  acid.  Of  the 
antiinoniates  may  be  mentioncHl  the  disodium  pyroautimoniate,  Na.dl^Sb./)^  -j  hi  1^0, 
Avhich  is  insoluble  in  cold  water,  and  is  therefore  precipitated  from  the  soluble  sodium 
salts  on  the  addition  of  dipotassium- pyroautimoniate. 

The  phosphates  show  more  plainly  than  the  sulphates  that  the  hydrogen  atoms  of  a 
polyhydric  acid,  replaceable  by  metals,  are  not  of  equal  importance  for  the  “strength  ” 
of  the  acid,  'fhe  hydrogen  first  replaced  acts  like  the  hydrogen  of  a strong  acid,  while 
the  second  and  third  follow  it  successively.  This  is  also  seen  in  the  reactions  of  the 
aqueous  solutions  and  may  be  ex|)lained  by  the  theory  of  electrolytic  dissociation  as 
follows:  'Phe  dissociation  of  phosphoric  acid  into  the  ions  and  11'  is  that  of 

an  acid  of  medium  strength  (p.  269).  'Phe  second  hydrogen  atom  is  dissociated  like  a 
feeble  acid  and  the  third  is  not  at  all  dissociated  in  aqueous  solution.  Consefjuently  the 
solution  of  sodium  triphosphate  does  not  contain  the  trivalent  anion  PO^'"  together  with 
the  sodium  ions.  As  the  third  hydrogen  atom  of  the  acid  shows  less  tendency  to  dis.so- 
ciate  than  water  the  followdng  transposition  immediately  takes  place  : PO^'"  -j- 

H'  HO'  P()^H"  -|-  HO',  so  that  the  solution  contains  the  ions  PO^H"  — 2Na' 

and  Na'  — OH',  which  are  the  cause  of  its  alkaline  reaction. 

Sodium  Borate. — The  normal  salts  of  boric  acid,  B(0H)3,  and 
metaboric  acid,  BO . OH  (see  p.  242),  are  not  very  stable.  The  ordi- 
nary alkaline  borates  are  derived  from  tetraboric  acid  (HgB^O^),  which 
results  from  the  condensation  of  four  molecules  of  the  normal  boric  acid  : 

4B(0H)3-5H,0  = H,BA. 

The  most  important  of  the  salts  is  borax,  which  crystallizes  at  ordinary 
temperatures  with  ten  molecules  of  water  in  large  monoclinic  prisms, 
Na^B^O^  -f-  10H2O.  Borax  occurs  naturally  in  some  lakes  of  Thibet, 
whence  it  was  formerly  imported  under  the  name  of  tinkaL  At  present, 
it  is  prepared  artificially  by  boiling  or  fusing  boric  acid  with  sodium  car- 
bonate. At  ordinary  temj)eratures,  the  crystals  dissolve  in  14  parts  of 
water,  at  100°  in  one-half  ])art;  the  solution  has  a feeble  alkaline  reac- 
tion. When  heated  to  70°  rhombohedra  crystallize  from.tbe  concen- 
trated solution,  and  have  the  composition  NaJl^O^  -|-  sH^O,  formerly 
known  as  octahedral  borax.  Both  salts  puff  up  when  heated,  lose  water, 
and  yield  a white,  porous  mass  {burned  borax^,  which  fuses  at  880°  to  a 
transparent  vitreous  mass  (Na.^B^O^).  In  fusion  this  dissolves  many  metal- 
lic oxides,  forming  transjiarent  glasses  {borax  beads),  which  frequently 
jjossess  characteristic  colors;  thus  co])per  salts  give  a blue  and  chromic 
oxide  gives  a green  glass.  Therefore,  borax  may  be  em])loyed  in  blow- 
lu'jje  tests  for  the  detection  of  certain  metals.  Upon  this  property  of 


LITHIUM.  295 

dissolving  metallic  oxides  depends  the  apidication  of  borax  for  the  fusion 
and  soldering  of  metals. 

Sodium  Silicate  (sodium  water-glass)  is  analogous  to  the  i)otassium 
salt,  and  is  most  readily  obtained  by  fusing  quartz  with  sodium  sulphate 
and  charcoal. 

The  sulphur  compounds  of  sodium  are  also  analogous  to  those  of  potas- 
sium. 

Sodium  Nitride,  NaNg  (sodium  azoimide),  is  the  sodium  salt  of  hydrazoic  acid, 
Ngil.  It  results  from  derivatives  of  azoimide,  by  neutralizing  the  free  acid  or  by  decom- 
posing the  ammonium  salt  with  caustic  soda.  W.  Wislicenus  prepared  it  by  heating 
sodamide  (p.  283)  to  150-250°  in  a current  of  nitrous  oxide  : 

NaNIIg  + N2O  = NaNg  -f  H,0. 

The  water  produced  here  decomposes  a portion  of  sodamide  into  sodium  hydroxide  and 
ammonia : 

NaNHg  + H2O  = NaOH  -f  NHg 

[Ber.  25  (1892),  2084;  see  also  Z.  f.  anorg.  Ch.  6 (1894),  38].  It  can  be  recrystallized 
from  water  or  may  be  precipitated  by  alcohol  from  water.  Its  solution  reacts  alkaline 
and  has  a very  salty  taste.  It  is  not  exploded  by  a blow,  but  this  occurs  when  it  is 
heated.  [Curtius,  Ber.  24  (1891),  3346;  see  also  p.  133.] 


Recognition  of  Sodium  Compounds. — Almost  all  the  sodium 
salts  are  easily  soluble  in  water,  sodium  pyroantimoniate,  Na2H2Sb20y, 
excepted;  this  is  precipitated  from  solutions  of  sodium  salts  by  potassium 
pyroantimoniate,  and  can  serve  for  the  detection  of  sodium.  Sodium 
compounds,  exposed  in  a colorless  flame,  impart  to  the  latter  an  intense 
yellow.  The  spectrum  of  the  sodium  flame  is  characterized  by  a very 
bright  yellow  line,  which,  when  more  strongly  magnified,  splits  into  two 
lines.  (See  Spectrum  Analysis.) 


4.  LITHIUM. 

Li  = 7.03. 

Lithium  occurs  in  nature  only  in  small  quantities,  but  is  tolerably 
widely  disseminated,  and  is  found  in  some  mineral  springs  and  in  the 
ashes  of  many  plants,  notably  in  that  of  tobacco  and  the  beet.  As  a com- 
pound silicate,  it  occurs  in  lepidolite  or  lithia  mica;  as  phosphate  (with 
iron  and  manganese)  in  triphylite,  and  (with  aluminium,  sodium,  and 
fluorine)  in  amblygonite. 

The  metal  is  separated  from  the  chloride,  or,  better,  from  the  more 
easily  fusible  mixture  of  equal  ])arts  of  lithium  chloride  and  potassium 
chloride,  by  means  of  the  electric  current,  and  is  silver-white  in  color, 
decomposing  water  at  ordinary  temperatures.  Its  specific  gravity  is  0.59. 
It  is  the  lightest  of  all  the  metals,  and  swims  upon  naphtha.  It  melts  at 
180°,  and  burns  with  an  intense  light.  It  burns  energetically  in  hydro- 
gen at  a red  heat  to  lithium  hydride,  Lill,  which  is  a comparatively  stable 


296 


INORGANIC  CHEMISI'RY. 


wliile  powder.  The  lilhiuiii  salts  arc  at  present  j)re|)ared  almost  entirely 
from  amblygonite  ; they  are  very  similar  to  the  salts  of  sodium,  but  closely 
approach  those  of  magnesium. 

Lithium  Chloride,  LiCd,  crystallizes,  at  ordinary  temperatures,  in 
anhydrous,  regular  octahedra  ; below  10°,  however,  it  has  two  molecules 
of  water,  and  delicpiesces  in  the  air. 

Lithium  Phosphate,  and  Lithium  Carbonate, 

Li^C(),j,  are  difficultly  soluble  in  water;  therefore  they  are  precipitated 
from  solutions  of  lithium  salts  by  sodium  j)hosi)hate  or  carbonate.  l>y 
strong  ignition  the  carbonate  loses  carbon  dioxide.  So  far  as  the.se  two 
salts  and  also  lithium  fluoride,  which  is  soluble  in  very  little  water,  are 
concerned,  lithium  aiijiroaches  the  metals  of  the  second  group  (|).  273). 
Its  compounds  color  the  flame  a beautiful  red;  the  spectrum  shows  an 
intense  red  line  together  with  a faint  yellow  line. 


AMMONIUM  COMPOUNDS. 

Upon  p.  128  we  observed  that  ammonia  combines  directly  with  the 
acids  to  form  salt-like  compounds,  which  are  analogous  to  the  metallic 
salts,  especially  those  of  potassium  with  which  they  are  isomorphous.  The 
univalent  group,  NH^,  playing  the  role  of  metal  in  these  derivatives,  is 
called  ammofiium,  and  the  derivatives  of  ammonia,  ammonium  coin- 
pounds.  The  metallic  character  of  the  group  NH^  is  confirmed  by  the 
existence  of  ammonium  amalgam,  which,  as  regards  its  external  appear- 
ance, is  very  similar  to  the  sodium  and  potassium  amalgams.  Ammonium 
amalgam  may  be  prepared  by  letting  the  electric  current  act  upon  ammo- 
nium chloride,  NH^Cl,  viz.,  by  immersing  the  negative  platinum  elec- 
trode into  a depression  in  the  ammonium  chloride,  which  is  filled  with 
mercury  and  stands  upon  the  positive  electrode.  Then,  as  in  the  case  of 
the  decomposition  of  potassium  or  sodium  chloride,  the  metallic  ion — 
ammonium — separates  on  the  negative  pole,  and  combines  to  an  amalgam 
with  mercury.  The  amalgam  may  also  be  obtained  if  sodium  amalgam 
be  covered  with  a concentrated  solution  of  ammonium  chloride: 

(Hg  4-  Na)  and  NH.d  yield  (Hg  + NH,)  and  NaCl. 

Sodium  amalgam.  Ammotiium  amalgam. 

Ammonium  amalgam  forms  a very  voluminous  mass  with  a metallic 
appearance.  It  is  very  unstable,  and  decomposes  rapidly  into  mercury, 
ammonia  and  hydrogen. 

The  ac|ueous  solution  of  ammonia  reacts  strongly  alkaline,  and  from  its 
entire  behavior  we  must  assume  the  existence  of  ammonium  hydroxide 
(NH/>M)  in  the  solution.  This  is  justified  because  there  are  many 
organic  derivatives  of  ammonium  hydroxide,  in  which  the  hydrogen  of 
the  ammonium  is  re|)laced  bv  hvdrocarbon  residues;  e.  g.,  tetramethyl 
ammonium  hydroxide,  N(CIi,,), OH.  'These  are  thick  liquids,  of  strong 
basic  reaction  and,  in  all  respects,  are  very  similar  to  potassium  and 
sodium  hydroxides. 


AMMONIUM  COMPOUNDS. 


297 


Ammonium  Chloride,  NH^Cl  {Sal ammoniac wyi),  is  sometimes  found 
in  volcanic  districts,  and  was  formerly  obtained  by  the  dry  distillation  of 
camel’s  dung  (p.  1 26).  At  present  it  is  prepared  almost  exclusively  from  the 
ammonia  water  of  gas  works.  This  water  contains  ammonium  sesquicar- 
bonate  in  addition  to  theotherammonium  salts.  It  is  distilled  with  lime 
and  the  escaping  ammonia  caught  in  hydrochloric  or  sulphuric  acid.  The 
ammonium  chloride  is  carefully  heated  and  recrystallized  or  sublimed.  By 
sublimation  it  is  obtained  as  a compact,  fibrous  mass.  It  dissolves  in  2.7 
parts  of  cold  and  in  one  part  of  boiling  water,  and  crystallizes  from  the 
solution  in  small,  feather-like,  grouped  octahedra  or  cubes,  of  sharp, 
salty  taste.  When  heated,  ammonium  chloride  sublimes  without  melting; 
at  the  same  time  a dissociation  into  ammonia  and  hydrochloric  acid  is 
sustained,  but  these  products  recombine  to  ammonium  chloride  on  cool- 
ing. The  dissociation  begins  at  280°  and  is  complete  at  350°,  and  the  vapor 
density  corresponds  to  that  of  a mixture  of  similar  molecules,  of  NHjand 
HCl,  i.  €.,  J7-07  f 3646_  _ 26.76  (02  = 32).  A like  decomposition  is  sus- 
tained by  the  ammonium  chloride  when  its  solution  is  boiled  ; ammonia 
esca])es  and  the  solution  contains  some  free  hydrochloric  acid. 

Ammonium  Sulphate,  (NHO2SO4,  crystallizes  without  water  in 
rhombic  prisms,  and  is  soluble  in  two  i)arts  of  cold  and  one  part  of  hot 
water.  It  fuses  at  140°,  and  by  further  heating  decomj^ioses.  Most  of  it 
is  used  as  a fertilizer. 

Ammonium  Persulphate  (pp.  188,  279),  (NH4)2S20,^,  obtained  by  the  electrolysis 
of  ammonium  sulphate,  is  at  present  made  on  a large  scale  and  is  used  as  an  oxidant.  It 
consists  of  very  soluble  monoclinic  crystals.  It  decomposes  when  its  aqueous  solution  is 
evaporated,  yielding  ammonium  sulphate,  free  sulphuric  acid  and  oxygen.  This  salt  is 
applied  in  making  the  other  persulphates. 

Ammonium  Nitrate,  NH^NOg,  is  isomorphous  with  potassium 
nitrate  and  deliquesces  in  the  air.  It  melts  at  159°  ; at  170°  decomposi- 
tion into  nitrous  oxide  and  water  commences  and  is  tumultuous  at  240° 
(p.  211).  It  has  been  recently  applied  in  the  manufacture  of  blasting 
material. 

Ammonium  Nitrite,  NH^N02,  is  present  in  minute  quantities  in  the 
air,  and  results  from  the  action  of  the  electric  spark  upon  the  latter  when 
moist.  It  may  be  obtained  by  the  saturation  of  aqueous  ammonia  with 
nitrous  acid  [Z.  f.  anorg.  Ch.  7 (1894),  34],  and  in  a perfectly  pure 
condition  by  the  decomposition  of  silver  or  lead  nitrite  by  ammonium 
chloride.  Heat  decomposes  it,  especially  when  in  concentrated  solution, 
into  nitrogen  and  water  (p.  115). 

The  decomposition  of  ammonium  nitrite  into  water  and  nitrogen  and  ammonium 
nitrate  into  nitrous  oxide  and  water  are  both  exothermic  reactions,  occurring  with  the 
disengagement  of  heat  and  are  independent  of  the  j^ressure  of  the  disengaged  gas  ; the 
components  do  not  reunite  to  form  their  original  compounds.  This  is  not  a case  of  dis- 
sociation fp.  274). 

Ammonium  Hyponitrite,  NIl4-0-N=N-0-NIl4,  has  been  pre]:)ared  recently 
by  conducting  ammonia  into  an  ethereal  solution  of  hy])onitrous  acid.  It  forms  white 
crystals.  It  melts  at  64°  with  violent  decomposition.  It  gradually  breaks  down  at  the 
ordinary  temperature  into  ammonia,  water  and  nitrous  oxide  [see  p.  212  and  Hantzsch 
and  Kaufmann,  Ann.  Chem.  292  (1896),  317]. 


298 


INORGANIC  CHF.MISTKY. 


Ammonium  Carbonate.— 'riu;  ncufralnx  secoiUaty (MI, )./;(),„ 
separates  as  a crystalline  powder,  when  aininonia  gas  is  condncied  ihrcnigli 
a concentrated  solution  of  commercial  ammonium  carbonate.  It  ])arts 
with  ammonia  in  the  air  and  becomes  the  primary  or  acid  salt,  NI  I,HC( >>3, 
which,  when  heated  to  58°,  dissociates  into  carbon  dioxide,  ammonia, 
and  water. 

d'he  common  commercial,  so-called  sesquicarbonate  of  ammonium,  is 
generally  a mixture  of  ])rimary  ammonium  carbonate  with  ammonium 
carbamate  (NHjIiCC).,  NH.^CCA^.  NH,.  d'he  latter  may  be  obtained 
by  the  direct  union  of  carbon  dioxide  with  ammonia;  water  immedi- 
ately converts  it  into  the  neutral  salt  (p.  234).  lOxjjerience  has  demon- 
strated that  the  commercial  salt  rpiite  often  contains  carbon  dioxide  and 
ammonia  in  the  same  proportion  as  the  acid  salt,  (Nlt  jHCO,.  It  arises 
in  the  decay  of  many  nitrogenous  carbon  comi)Oiinds,  e.  g.,  the  urine, 
and  was  formerly  prepared  by  the  dry  distillation  of  bones,  horn,  and 
other  animal  substances.  At  ])resent  it  is  obtained  by  heating  a mixture 
of  ammonium  chloride,  or  sulphate,  with  calcium  carbonate.  It  then 
sublimes  as  a white,  tr;ins])arent,  hard  mass. 

Primary  Ammonium  Carbonate,  NH^HCO.,,  is  obtained  by  satu- 
rating ammonium  hydroxide  with  carbon  dioxide.  It  is  a white,  odor- 
less powder,  rather  insoluble  in  water.  In  aqueous  solution  it  gradually 
loses  carbon  dioxide  and  is  changed  to  the  secondary  carbonate. 

Ammonium  Phosphates. — 'Fhe  most  important  of  these  is  the 
secondary  ammo7iium-sodium  phosphate,  NH^NaHFO^  -f-  qll^O,  ordi- 
narily termed  salt  of  phosphorus  (Sal  ?nicrocosmicum').  It  is  found  in 
guano  and  in  decaying  urine.  It  can  be  obtained  by  the  crystallization 
of  a mixture  of  disodium  phosphate  and  ammonium  chloride  : 

Na,HPO,  -f  NH,C1  = NI^NallPO,  + NaCl. 

It  consists  of  large,  transparent,  monoclinic  crystals.  When  heated  it 
fuses,  giving  up  water  and  ammonia  and  forms  a transparent  glass  of 
sodium  metaphosphate,  NaPOg  (p.  294).  It  serves  in  blowpipe  tests  for 
the  detection  of  various  metals. 

The  tertiary  ammonium  phosphate,  (NHJ3P0^,  separates,  in  crystalline 
form,  upon  mixing  concentrated  solutions  of  phosphoric  acid  and  ammo- 
nia. Upon  drying,  it  loses  ammonia  and  passes  into  the  secondary  salt, 
(NHJ.^HPO^,  which  changes  to  the  prhnary  salt,  (NH^)H2PO^,  when  its 
solution  is  boiled.  This  is  in  harmony  with  what  was  said  on  p.  294  rela- 
tive to  the  behavior  of  j)hosi)horic  acid. 

Ammonium  Nitride,  N./NIip,  the  ammonium  salt  of  hydrazoic  acid,  obtained  from 
an  organic  body,  diazohi])])uramide,  or  by  saturating  hydrazoic  acid  vvitli  ammonia,  is  pre- 
ci])itated  from  its  alcoliolic  solution  byetlier  in  the  form  of  a snow-white,  crystalline  pow- 
der. It  separates  from  alcoliol  in  com|)act,  colorless  leaflets,  consisting  of  step-like 
groups  of  crystals.  It  resembles  ammonium  chloride  in  this  res])ect.  It  crystallizes  from 
water  in  large,  (rans|\arent  ])risms,  which  soon  become  opaque.  It  has  a slight  alkaline 
reaction  ; it  is  not  hygroscopic  ; it  dissolves  readily  in  water  and  in  alcohol.  It  is 
exceedingly  volatile  ; it  graduallv  disappears  in  the  air  and  it  is  also  carried  off  by 
af)ueous  and  alcoholic  vapors.  It  sublimes,  when  gently  heated,  in  small,  shining 
prisms;  on  rapid  healing  it  explodes  (Clurtius,  [).  132). 


METALS  OF  THE  SECOND  GROUP. 


299 

Ammonium  Sulphide,  (NHJ,^S,  results  ii[)on  mixing  i volume  of 
hydrogen  sulphide  with  2 volumes  of  ammonia  at  — 18°.  It  is  a white 
crystalline  mass,  dissociating,  at  ordinary  temperatures,  into  NH^SH  and 
NHg.  In  aqueous  solution  it  also  seems  to  dissociate  into  its  constituents. 
At  45°  it  completely  dissociates  into  ammonia  and  hydrogen  sulphide ; 

(NHJ2S  = 2NII3  + H^S. 

2 vols.  I vol. 

Ammonium  Hydrosulphide,  NH^SH,  is  produced  upon  conduct- 
ing hydrogen  sulphide  into  an  alcoholic  ammonia  solution.  It  is  com- 
pletely dissociated  at  45°  : 

NH^SH  = NHg  -f  HgS. 

I vol.  1 vol. 

It  is  obtained  in  aqueous  solution  by  saturating  aqueous  ammonia  with 
hydrogen  sulphide.  At  first  the  solution  is  colorless,  but  on  standing  in 
contact  with  the  air  becomes  yellow,  owing  to  the  formation  of  ammonium 
polysulphides,  (NH^}2Sn.  The  so-called  yellow  ammonium  sulphide  is 
more  easily  obtained  by  the  solution  of  sulphur  in  the  colorless  hydrosul- 
phide. Both  solutions  are  often  employed  in  laboratories  for  analytical 
purposes. 

Recognition  of  Ammonium  Compounds. — All  ammonium  salts 
are  volatile  or  decompose  upon  heating.  The  alkalies  and  other  bases 
liberate  ammonia  from  them,  which  is  recognized  by  its  odor  and  the 
blue  color  it  imparts  to  red  litmus-paper.  Platinum  chloride  produces 
a yellow  crystalline  precipitate  of  ainmonio-platinum  chloride,  PtCl^.- 
2NH^C1,  in  solutions  of  ammonium  chloride.  An  excess  of  tartaric  acid 
precipitates  primary  ammonium  tartrate. 


METALS  OF  THE  SECOND  GROUP. 


Be  9. 1 


Mg  24.36 


Ca  40  Sr  87.6  Ba  137.4 


Zn  65.4  Cd  112 


Ilg  200.3. 


The  second  group  of  the  periodic  system  (see  Table,  p 246)  comprises 
chiefly  the  bivalent  metals,  which  form  compounds  of  the  type  MeXg,  and 
in  their  entire  deportment  exhibit  many  analogies.  Their  special  rela- 
tions and  analogies  are  more  closely  regulated  by  their  position  in  the 
periodic  system.  Beryllium  and  magnesium  belong  to  the  two  small 
periods  whose  members  are  similar  but  do  not  show  complete  analogy. 
Beryllium  exhibits  many  variations  from  magnesium,  and  in  manv  prop- 
erties approaches  aluminium;  just  as  lithium  attaches  itself  to  magnesium 
(P;  273)- 

The  metals  calcium,  strontium,  and  barium  constitute  the  second 
members  of  the  three  great  periods,  are  very  similar  to  one  another  (p 
244),  and  in  accord  with  their  strong  basic  character,  attach  themselves 
to  the  alkali  metals — potassium,  rubidium  and  ctesium.  Zinc,  cadmium. 


300 


INORGANIC  CHEMISTRY. 


and  mercury,  wliicli  correspond  to  them  and  constitute  the  second  sub- 
group, really  belong  to  the  right,  negative  sides  of  the  three  great  jjeriods. 
'They  fall  in  with  the  heavy  metals,  are  much  less  basic,  and  resemble  the 
alkaline  earth  metals  only  in  their  combination  types.  In  conseipience 
of  the  double  periodicity  of  the  three  great  periods  both  sub-groui)S  (Ca, 
Sr,  Ba  and  Zn,  Cd,  Hg)  exhibit  many  analogies  to  magnesium  and 
beryllium. 


I.  GROUP  OF  THE  ALKALINE  EARTHS. 

Calcium,  . . Ca  = 40  Strontium,  . . Sr  = 87.6  IJarium,  , . P>a  = 137.4. 

The  metals  of  this  group  are  termed  alkaline  earth  metals,  because 
their  oxides  attach  themselves  in  their  i)roperties,  on  the  one  side  to  the 
oxides  of  the  alkalies,  upon  the  other  to  the  real  earths  (alumina,  etc.). 
'I'heir  atomic  weights  bear  almost  the  same  ratio  to  one  another  as  those 
of  the  alkali  metals,  hence  the  alkaline  earth  metals  show  the  same  grada- 
tion in  properties  as  the  elements  of  the  ])otassium  group.  With  in- 
crease in  atomic  weight  and  atomic  volume,  their  chemical  energy  and 
basicity  become  greater.  Barium  decomposes  water  energetically,  and 
oxidizes  more  readily  than  strontium  and  calcium.  In  accord  with  this, 
we  find  barium  hydroxide  a stronger  l)ase  ; it  dissolves  rather  readily  in 
water,  does  not  decompose  upon  ignition,  and  absorbs  carbon  dioxide 
rapidly  from  the  air.  Barium  carbonate  is  also  very  stable,  fuses  at  a 
white  heat,  and  only  disengages  a little  carbon  dioxide.  Calcium  hy- 
droxide, on  the  other  hand,  dissolves  with  more  difficulty  in  water,  and 
when  ignited,  breaks  down  into  water  and  calcium  oxide;  the  carbonate 
also  yields  carbon  dioxide  when  similarly  treated.  In  its  entire  charac- 
ter, strontium  stands  between  barium  and  calcium.  All  these  affinity  re- 
lations find  full  expression  in  the  heat  of  formation  of  the  corresponding 
compounds  (p.  327). 

While  the  alkaline  earth  metals  are  similar  to  the  alkalies  in  their 
free  condition  and  in  their  hydroxides,  they  differ  essentially  from  them 
by  the  insolubility  of  their  carbonates  and  phosphates,  and  still  more  of 
their  sulphates.  Barium  sulphate  is  only  soluble  to  the  slightest  degree 
in  water  and  acids,  while  strontium  sulphate  is  about  40  times  and  cal- 
cium sulphate  about  800  times  as  soluble  (p.  309). 

The  atomic  weights  of  the.se  metals  have  been  determined  in  part  from  their  specific 
lieats,  but  mainly  from  their  isomorphism  with  the  metals  of  the  magnesium  group 
(pp.  253,  255). 


1.  CALCIUM. 

Ca  = 40. 

Calcium  belongs  to  the  class  of  elements  most  widely  distributed  upon 
the  earth’s  surface.  As  calcium  carbonate  (limestone,  marble,  chalk) 
and  the  sulphate  (gyi)sum,  alabaster),  it  represents  immense  deposits  in 


CALCIUM. 


301 


all  stratified  formations.  As  phosphate,  it  constitutes  phosphorite,  as 
fluorite,  both  of  which  are  abundant.  As  silicate,  it  is  found  in 
most  of  the  oldest  crystalline  rocks. 

The  metal  is  obtained,  according  to  Moissan,  by  heating  calcium 
iodide  with  an  excess  of  sodium  to  a dark-red  heat.  The  liberated 
metal  dissolves  in  the  sodium,  and  when  cold  the  latter  is  removed  by 
anhydrous  alcohol.  It  dissolves  with  the  evolution  of  hydrogen  to 
sodium  alcoholate.  The  calcium  remains  as  a brilliant  white  crystalline 
powder.  It  can  also  be  prepared  by  the  electrolysis  of  the  fused  chloride 
or  iodide.  Although  the  affinity  of  calcium  for  oxygen  is  less  than  that 
of  the  alkalies,  yet  the  oxide  (also  barium  oxide  and  strontium  oxide) 
cannot  be  reduced  to  metal  by  ignition  with  carbon,  iron,  or  sodium — 
due,  probably  to  the  non-fusibility  of  the  oxide.  These  oxides,  how- 
ever, are  reduced  in  the  electric  furnace;  the  liberated  metals  com- 
bine at  once  with  carbon  to  yield  metallic  carbides  (pp.  253,  307). 

Calcium  is  a yellow,  shining  metal,  of  specific  gravity  1.55-1.6.  In 
dry  air  it  is  tolerably  stable,  in  moist  it  is  covered  with  a layer  of 
hydroxide.  It  decomposes  water  with  considerable  energy.  It  fuses  at 
a red  heat,  and  in  the  air  burns  with  a brilliant  yellow  light.  If  the 
metal  be  heated  in  nitrogen  to  a red  heat  calcium  nitride,  Ca3N2,  results. 
This  is  a brown  mass  which  water  decomposes  with  the  formation  of 
calcium  hydrate  and  ammonia.  Lithium,  strontium,  barium  and  mag- 
nesium behave  similarly  (p.  316).  Calcium  slowly  unites  with  hydrogen 
at  the  ordinary  temperature,  but  rapidly  at  a red  heat,  forming  calcium 
hydride,  CaHj,  an  earthy,  gray  powder,  which  decomposes  water  more 
energetically  than  calcium  itself. 

Calcium  Oxide,  CaO  (lime),  may  be  obtained  ])ure  by  igniting  the 
nitrate  or  carbonate.  It  is  prepared  on  a large  scale  by  burning  the 
ordinary  limestone  or  marble  (CaCOg)  in  lime-kilns.  It  is  a grayish- 
white  amorphous  mass,  which  becomes  crystalline  at  2500°  and  liquid 
like  water  at  3000°.  It  vaporizes  at  higher  temperatures  (Moissan). 
The  oxyhydrogen  flame  thrown  upon  a piece  of  lime  causes  it  to  emit  an 
extremely  intense  white  light  (Drummond’s  lime  light).  In  the  air 
lime  attracts  moisture  and  carbon  dioxide,  becoming  calcium  carbonate  ; 
burned  lime  unites  with  water  with  evolution  of  much  heat,  breaking 
down  into  a white  voluminous  powder  of  calcium  hydroxide,  Ca(OH)2 — 
slaked  Ihjie. 

When  limestone  contains  large  quantities  of  alumina,  magnesium  carbonate,  or  other 
constituents,  the  lime  from  it  slakes  with  difficulty,  and  is  known  as  poor  lime,  to  dis- 
tinguish it  from  pure  fat  or  rich  lime,  which  is  readily  converted  into  a powder  of 
hydroxide  with  water. 

Calcium  Hydroxide,  Ca(OH)2  (slaked  lime),  is  a white,  porous 
powder  which  dissolves  with  difficulty  in  cold  water  (i  part  in  760  parts), 
but  with  still  more  difficulty  in  warm  water;  the  solution  saturated  in 
the  cold  (lime-water)  becomes  cloudy  uj)on  warming.  It  has  a strong 
alkaline  reaction.  Milk  of  Imie  is  slaked  lime  mixed  with  water.  In  the 
air  it  attracts  carbon  dioxide  and  forms  calcium  carbonate.  At  a red 
heat  it  decomposes  into  oxide  and  water. 


302 


INORGANIC  CHEMISTRY. 


Slaked  lime  is  employed  in  the  j)re}jaration  of  ordinary  mortar,  a mix- 
ture of  calcium  hydroxide,  water  and  sand.  The  liardening  of  the  mor- 
tar in  the  air  depends  mainly  upon  the  fact  that  the  calcium  hydroxide 
combines  with  the  carbon  dioxide  of  the  air  to  form  the  carbonate,  and 
at  the  same  time  acts  upon  the  silicic  acid  of  the  sand  forming  a calcium 
silicate,  which,  in  time,  imi)aits  durability  to  the  mortar. 

Hydraulic  mortar,  or  cement,  is  produced  by  gently  igniting  a mixture 
of  limestone  or  chalk  with  aluminium  silicate  (clay)  and  (piartz  powder. 
On  stirring  the  jiowdered,  burnt  mass  with  water  it  soon  hardens,  and  is 
then  not  affected  by  water.  Some  naturally  occurring  limestones,  con- 
taining upwards  of  20  jier  cent,  of  clay,  yield  hydraulic  cements,  without 
any  admixtures  after  burning.  Their  comi)Osition  varies,  and  also  the 
])rocess  of  their  hardening  ; the  latter,  however,  depends  princij)al]y  upon 
the  formation  of  calcium  and  aluminium  silicates. 

Calcium  Peroxide,  Ca02,  is  precipitated  as  a hydrate  in  crystalline 
leaflets,  if  lime-water  be  added  to  a solution  of  hydrogen  jjeroxide.  It 
contains  eight  molecules  of  water,  which  it  gradually  loses  in  dry  air. 


The  halogen  derivatives  of  calcium,  like  those  of  other  metals,  are 
})repared  by  the  solution  of  the  oxide  or  carbonate  in  the  aqueous  haloid 
acids.  They  are  formed  also  by  the  direct  union  of  calcium  with  the 
halogens;  this  occurs  with  evolution  of  flame.  Technically,  calcium 
chloride  is  often  obtained  as  a by-product,  e.  g.,  in  the  preparation  of 
ammonia  (see  Soda). 

Calcium  Chloride,  CaCl2,  crystallizes  from  aqueous  solution  with  six 
molecules  of  water,  in  large,  transparent,  six-sided  prisms,  which  deli- 
quesce in  the  air.  In  vacuo  it  loses  four  molecules  of  water.  When 
heated,  it  melts  in  its  water  of  crystallization,  loses  water,  but  it  is  only 
after  it  has  been  exposed  above  200°  that  it  becomes  anhydrous;  then  it 
is  a white,  porous  mass.  The  dry  salt  melts  at  806°,  and  solidifies  on 
cooling  to  a crystalline  mass,  which  attracts  water  energetically,  and  may 
be  employed  in  the  drying  of  gases  and  liquids.  The  dry  calcium  chlo- 
ride also  absorbs  ammonia,  forming  the  compound  CaCl2 . 8NH3.  The 
crystallized  hydrous  salt  dissolves  in  water  with  reduction  of  temperature; 
by  mixing  with  snow  or  ice  the  temperature  is  lowered  to  ^ — ^_48°.  Upon 
fusing  the  dry  chloride  in  moist  air  it  will  partly  decompose  into  the 
oxide  and  hydrogen  chloride. 

Calcium  bromide  and  iodide  are  very  similar  to  the  chloride. 

Calcium  Fluoride,  CaFl.^,  occurs  in  nature  as  fluorite,  in  large  cubes 
or  octahedra,  or  even  in  compact  masses.  It  is  often  discolored  by  im- 
])urities.  It  is  found,  in  sparing  quantities,  in  the  ashes  of  idants,  bones, 
and  the  enamel  of  the  teeth.  A soluble  fluoride  added  to  the  solution  of 
calcium  chloride  throws  down  insoluble  calcium  fluoride  as  a white  vol- 
uminous ])recii)itate.  If,  however,  the  calcium  cldoride  solution  be  poured 
into  a boiling  dilute  solution  of  potassium  fluoride,  the  calcium  fluoride 
is  obtained  in  little  transparent  crystals. 

'I'he  fluoride  is  jierfectly  insoluble  in  water,  and  is  only  decomposed  by 


CALCIUM  HYPOCHLORITE. 


303 


Strong  acids.  It  fuses  easily  at  a red  heat,  serving,  therefore,  as  a flux  in 
the  smelting  of  ores.  When  heated  it  phosphoresces. 


Calcium  Hypochlorite,  Ca(C10)2,  is  not  known  in  a pure  condi- 
tion. The  so-called  bleaching  lime  or  chloride  of  lime,  obtained  by  con- 
ducting chlorine,  at  ordinary  temperatures,  over  slaked  lime,  contains 
calcium  hypochlorite  as  the  active  principle. 

According  to  analogy  to  the  action  of  chlorine  upon  potassium  or 
sodium  hydroxide  (p.  278),  the  reaction  in  the  case  of  calcium  hydroxide 
may  be  expressed  by  the  following  equation  : 


2Ca(OH)2  + 2CI2  = Ca(C10)2  -f-  CaCq  -f  2H2O. 


This  would  incline  us  to  regard  chloride  of  lime  as  a mixture  of  cal- 
cium hypochlorite,  calcium  chloride,  and  water.  In  accordance  with 
the  equation,  the  completely  chlorinated  chloride  of  lime  must  contain 
48.9  per  cent,  of  chlorine,  which  is  never  the  case,  because  a portion  of 
the  calcium  hydroxide  invariably  remains  unaltered.  Calcium  chloride 
does  not  exist  free  in  bleaching  lime,  because  it  is  not  withdrawn  from 
the  latter  by  alcohol,  and  nearly  all  the  chlorine  of  the  bleaching  lime 
can  be  expelled  by  carbon  dioxide.  It  is  therefore  probable  that  the 
Cl 

compound,  Ca<Q^I  = CaCl20  (Odlingand  Lunge),  is  present  in  bleach- 
ing lime;  its  formation,  then,  from  calcium  hydroxide  would  correspond 
to  that  of  the  chloride  and  hypochlorite  from  the  caustic  alkalies; 


NaOH  , 
NaOH  “t-  '-h 


Chloride  of  lime  is  a white,  porous  powder  with  an  odor  resembling 
that  of  chlorine.  The  aqueous  solution  has  a strong  alkaline  reaction, 
and  bleaches.  It  decomposes  in  the  air  as  the  carbon  dioxide  of  the 
latter  liberates  hypochlorous  acid.  Even  in  closed  vessels  it  gradually 
decomposes  with  the  evolution  of  oxygen  ; the  decomposition  is  hastened 
I)y  sunlight  and  heat,  and  may  occur  with  explosion.  Hence  chloride 
of  lime  should  be  preserved  in  loosely  closed  vessels,  in  a cool,  dark 
place. 

Dilute  hydrochloric  or  sulphuric  acid  will  expel  chlorine  from  chloride 
of  lime  ; the  quantity  liberated  is  just  twice  that  which  the  hypochlorite 
eventually  found  in  it  contains  (p.  175)  : 


CaCl(Ori)  -I-  2IICI  = CaCq  + 11,0  -f  Cl,. 


By  the  action  of  sulphuric  acid  all  the  chlorine  of  the  bleaching 
powder  is  liberated,  whereby  calcium  chloride  first  yields  hydrochloric 
acid : 


CaCl(OCl)  + = CaSO^  + 11,0  f Cl,. 


304 


INORGANIC  CHEMISTRY. 


The  api)lication  of  chloride  of  lime  for  the  production  of  chlorine  in 
chlorine  bleaching  and  disinfection  is  based  on  this  deiiorlment. 

The  (juantity  of  chlorine  set  free  by  acids  from  the  chloride  of  lime 
rei)resents  its  quantity  of  so-called  active  chlorine ; good  chloride  of  lime 
should  contain  at  least  25  per  cent. 

When  a small  (iuantity  of  coballic  oxide  or  a cobalt  salt  is  added  to  the  solution  of 
bleaching  lime,  and  heat  aj)[)lied,  a regular  stream  of  oxygen  is  disengaged  ; this  is  an 
advantageous  method  of  preparing  oxygen.  Other  oxides,  like  those  of  manganese, 
copper,  and  iron,  behave  similarly.  In  this  reaction  there  occurs,  apparently,  a cata- 
lytic action  of  the  oxides,  d’he  reaction  is  explained,  doubtless,  in  the  same  way  as  the 
action  of  hydrogen  peroxide  upon  certain  oxides  (sec  pp.  8S,  102). 

On  warming  bleaching  lime  with  ammonia,  the  following  decomposition  occurs  : 

3CaCl(OCl)  I 2NII3  = 3CaCb  + 3H,0  + N^. 

This  reaction  can  be  used  for  the  preparation  of  nitrogen. 

Calcium  Sulphate,  CaSO^,  is  very  abundant  in  nature.  In  an 
anhydrous  condition  it  forms  the  mineral  anhydrite,  crystallizing  in  forms 
of  the  rhombic  system.  With  two  molectiles  of  water  it  occurs  as 
gypsum,  in  large  monoclinic  crystals  or  in  granular,  crystalline  masses 
(alabaster,  etc.).  It  also  separates  as  a fine  crystalline  ])owder,  CaSO^  d- 
2H2O,  vv’hen  soluble  calcium  salts  are  precipitated  with  sulphuric  acid. 
Calcium  sulphate  dissolves  with  difficulty  in  water;  i ])art  at  average 
temperatures  is  soluble  in  400  parts  of  water.  When  heated  to  110° 
gypsum  loses  all  its  water,  and  becomes  burjtt  g^\)’s,\\m  ; when  this  is  pul- 
verized and  mixed  with  water,  it  forms  a paste  which  hardens  to  a solid 
mass  in  a short  time.  The  hardening  is  dependent  ui)on  the  reunion  of 
anhydrous  calcium  sulphate  with  two  molecules  of  water.  On  this  de- 
pends the  use  of  burned  gypsum  for  the  production  of  casts,  figures, 
etc.  In  case  gypsum  is  heated  above  160°  (dead-burnt  gypsum),  it  will 
no  longer  harden  with  water  ; the  naturally  occurring  anhydrite  behaves 
in  the  same  manner. 

Calcium  Nitrate,  Ca(N03)2,  is  produced  by  the  decay  of  the  nitrog- 
enous organic  substances  in  the  presence  of  lime,  therefore,  it  is  fre- 
quently found  as  an  efflorescence  upon  walls  (in  cattle  stables).  It 
crystallizes  from  water  in  monoclinic  prisms,  containing  four  molecules 
of  water;  the  anhydrous  salt  deliquesces  in  the  air  and  is  easily  soluble 
in  alcohol.  By  the  action  of  potassium  carbonate  or  chloride,  calcium 
nitrate  can  be  transposed  into  ])otassium  nitrate  (p.  280). 

Calcium  Phosphates. — The  tertiary  phosphate,  Ca3(PO^)2,  is  found 
in  slight  quantities  in  most  of  the  mountain  rocks.  In  combination  with 
calcium  chloride  and  fluoride,  it  crystallizes  as  apatite.  As  phosphorite, 
it  forms  compact  masses,  more  or  less  intimately  mixed  with  other  con- 
stituents, and  occurs  in  immense  deposits  in  Spain,  France,  Germany, 
Russia  and  Florida.  When  these  minerals  disintegrate  the  calcium  phos- 
])hate  ])asses  into  the  soil  and  is  absorbed  by  the  plants.  In  the  latter, 
it  accumulates  chiefly  in  the  seeds  and  grains.  In  the  animal  kingdom, 
it  is  principally  found  in  the  bones,  the  ashes  of  which  contain  upwards 
of  82  per  cent,  of  tricalcium  ])hosj)hate.  T'ertiary  calcium  phosphate 
is  nearly  insoluble  in  water.  If  disodium  phosphate  be  added  to  the 


CALCIUM  CARBONATE. 


305 


aqueous  solution  of  a calcium  salt,  and  then  ammonium  hydroxide,  it 
will  separate  as  a gelatinous  precipitate,  which,  after  drying,  becomes  a 
white  amorphous  powder.  It  is  very  readily  soluble  in  acids,  even  acetic. 

The  secondary  calcium  phosphate,  CaHPO^  -f  2H2O,  is  sometimes 
present  in  guano,  in  the  form  of  small,  shining  prisms;  it  separates  as 
an.amorphous  precipitate,  if  disodium  phosphate  be  added  to  a solution 
of  calcium  chloride  mixed  with  some  acetic  acid.  When  ignited,  it 
passes  into  pyrophosphate,  Ca.^P.^O.^  (p.  216). 

The  primary  phosphate,  Ca(H.,P04).„  is  produced  by  the  action  of  sul- 
phuric or  hydrochloric  acid  upon  the  first  two  phosphates.  It  is  readily 
soluble  in  water.  It  changes  to  metaphosphate  when  it  is  heated  : 

Ca(ITPOj2  - Ca(P03b  + 211,0. 

Calcium  phosphate  is  present  in  all  plants.  Its  presence  in  the  soil  is,  therefore,  an 
indispensable  condition  for  its  fertility.  When  there  is  a scarcity  of  phosphoric  acid  it 
must  be  added.  To  this  end,  bone  meal  and  pulverized  phosphorite  were  formerly  em- 
ployed. Since,  however,  the  phosphoric  acid  is  contained  in  these  substances  as  tri- 
calcium phosphate,  which  is  not  easily  absorbed  by  the  plants,  the  primary  phosphate  is 
extensively  employed  at  present  as  a fertilizer,  or,  better,  the  mixture  resulting  from  the 
action  of  sulphuric  acid  upon  the  tertiary  salt,  which  also  contains  calcium  sulphate. 
Superphosphate  is  the  name  applied  to  the  resulting  mass. 

The  Thomas  slag,  obtained  in  the  dephosphorization  of  iron  oi'es  by  the  Gilchrist- 
Thomas  process,  constitutes  a very  valuable  and  important  crude  product  from  which  to 
prepare  calcium  phosphate.  It  also  contains  a crystallized  phosphate  of  the  formula 
Ca^F.^O;,,  which  may  be  regarded  as  a derivative  of  the  normal  hydrate,  r(0H)5  : 

2P(0H)5—  H.,0  = HgP^Og. 

Calcium  Carbonate,  CaCOg,  is  very  widely  distributed  in  nature. 
It  crystallizes  in  two  crystallographic  systems,  hence  is  dimorphous.  In 
rhombic  crystals,  with  the  specific  gravity  3.0,  it  forms  aragonite.  In 
hexagonal  rhombohedra,  with  the  specific  gravity  2. 7,  it  occurs  as  calcite. 
Iceland  spar,  employed  for  optical  purposes,  is  perfectly  pure,  transparent 
calcite.  The  common  calcite,  which  constitutes  immense  mountain 
chains,  is  an  amorphous  or  indistinctly  crystalline  stratum,  and  is  usually 
mixed  with  other  constituents,  as  clay.  When  the  limestone  is  granular 
and  crystalline,  it  is  termed  marble.  Dolomite  also  forms  large  layers, 
and  is  a compound  or  amorphous  mixture  of  calcium  and  magnesium 
carbonates,  with  generally  an  excess  of  the  former.  Chalk  is  very  pure 
amorphous  calcium  carbonate,  consisting  of  the  shells  of  microscopic  sea 
animals.  Calcium  carbonate  is,  further,  a regular  constituent  of  all  plants 
and  animals ; the  shells  of  eggs,  of  mussels,  even  corals  and  pearls,  consist 
chiefly  of  it. 

A soluble  carbonate,  added  to  the  aqueous  solution  of  a calcium  salt, 
precipitates  calcium  carbonate  as  a white,  amorphous  powder,  which  soon 
becomes  crystalline.  In  the  cold,  it  assumes  the  form  of  calcite  ; upon 
boiling  the  liquid,  it  generally  changes  into  aragonite  crystals.  Magnesite 
(MgCOg),  siderite  (P'eCO.j),  smithsonite  (ZnCO.^)  and  rhodochrosite 
(MnCOg)  are  isomorphous  with  calcite,  while  witherite  (BaCO,,),  stron- 
tianite  (SrCOg)  and  cerinssite  (PbC03)are  isomorphous  with  aragonite. 

Calcium  carbonate  is  almost  insoluble  in  pure  water;  but  dissolves 
somewhat  in  water  containing  carbon  dioxide,  as  it  very  probably  is 
26 


3o6 


INORGANIC  CHEMISTRY. 


changed  lo  ihe  primary  carljonate,  Ca(HCC).,)2.  One  tliousand  parts  of 
water  at  15°  and  incdiuin  barometric  j)res.siire  dissolve  about  0.385  part 
of  ])rimary  calcium  carbonate  [Z.  f.  anorg.  Ch.  17  (1898)  170].  Id)r 
this  reason,  we  find  calcium  bicarbonate  dissolved  in  all  natural  waters. 
When  the  solution  stands  ex|)Osed,  or  if  it  be  heated,  carbon  dioxide 
escai)es,  and  the  secondary  carbonate  again  separates  out.  The  formation 
of  lime  scales,  thermal  tufts,  stalactites,  boiler  scales  and  similar  deposits 
are  due  to  this.  Calcium  carbonate,  like  all  carbonates,  is  decomj)osed 
by  acids,  with  evolution  of  carbon  dioxide. 

At  a red  heat,  it  dissociates  into  calcium  oxide  and  carbon  dioxide.  Tlie  change 
begins  at  600°.  Tlie  tension  of  dissociation  (]).  274)  is  85  mm.  at  860°,  and  at  1040° 
510-520111111.  For  this  reason  calcium  carbonate  is  not  decomposed  wlien  heated  in  a 
sealed  tube. 

Calcium  Silicate,  CaSiO,,  occurs  as  white,  crystalline  wollastonite. 
It  is  also  a constituent  of  most  natural  silicates  and  of  the  artificial  sili- 
cate fusion — glass. 

Glass. — The  silicates  of  pota.ssinm  and  sodium  are  readily  fusible  and  soluble  in  water. 
The  silicates  of  calcium  and  the  other  alkaline  earths  are  insoluble,  very  difficultly  fusible, 
and  generally  crystallize  when  they  cool.  If,  however,  the  two  silicates  be  fused  together, 
ail  amorphous,  transparent  mass,  of  average  fusibility,  results  ; it  is  only  slightly  attacked 
by  water  and  acids — it  is  glass.  To  prejiare  the  latter,  a mixture  of  sand,  lime,  and 
soda,  or  potash,  is  heated  to  fusion  in  a muffle  furnace.  Instead  of  the  carbonates  of 
potassium  and  sodium  a mixture  of  the  sulphates  with  charcoal  can  be  employed  ; the  carbon 
reduces  the  sulphates  to  sulphides,  which  form  silicates  when  fused  with  silicon  dioxide. 

Glass,  depending  on  its  constituents,  arranges  itself  into  two  clas.ses  : 

1.  Lime  glass  (free  from  lead). 

2.  Lead  glass. 

1.  The  lime  glass eonsists either  of  sodium-lime  (soda  glass)  or  potassium-lime  (potash 
glass). 

Soda  g/ass — a mixture  of  sodium  and  calcium  silicates — buses  readily,  and  is  employed 
for  window-panes  and  ordinary  glass  vessels.  Poiash  or  BoJnviian  glass,  also  called 
crown  glass,  consists  of  calcium  and  potassium  silicates,  is  not  very  fusible,  is  harder 
and  withstands  the  action  of  water  and  acids  better  than  soda  glass  ; it  is,  therefore, 
employed  in  the  manufacture  of  chemical  glassware. 

The  best  lime  glass,  normal  glass,  is  a trisilicate  corresponding  to  the  formula- 

KgO  (or  Na20)  CaO  -f-  bSiOg. 

2.  Lead  glass  is  constituted  like  lime  glass,  eontaining,  however,  lead  oxide  for 
calcium  oxide. 

Glass  crystal  or  flint  glass  is  composed  of  potassium  and  lead  silicates.  It  is  not  as 
hard,  fuses  with  tolerable  readiness,  refracts  light  strongly,  and  when  polished,  acquires 
a clear  luster.  On  this  account  it  is  employed  for  optical  purposes  (for  lenses,  prisms) 
and  is  u.sed  in  ornamental  glassware.  Strass,  a lead  glass  containing  boron  trioxide,  is 
used  to  imitate  ])recious  stones.  The  opaque  varieties  of  enamel  consist  of  lead  glass 
and  in  the  fused  glass  are  insoluble  admixtures,  as  tin  dioxide  and  calcium  phosphate. 

Ordinary  winclow  or  bottle  glass  is  obtained  by  the  fusion  of  rather  impure  materials  ; 
in  consc(]uence  of  Ihe  ])resenceof  ferrous  oxide  it  is  ordinarily  colored  green.  To  remove 
this  coloration,  manganese  peroxide,  MnO,^,  is  added  to  the  fusion.  It  oxidizes  a por- 
tion of  the  ferrous  to  ferric  oxide,  the  silicate  of  which  is  colored  slightly  yellow,  while 
manganese  foi  ins  a violet  silicate.  'I'liese  colors,  violet  and  green,  almost  neutralize  each 
other  as  complementaries.  d’he  colored  glasses  contain  silicates  of  colored  metallic 
oxides:  chromic  and  cn|)ric  oxides  color  green;  cobaltic  oxide,  blue  ; cupn)us  oxide,  a 
ruby  red,  etc. 

'The  sulphur  compounds  of  calcium  are  very  much  like  those  of  the  alkalies.  Calcium 


STRONTIUM. 


307 


Sulphide,  CaS,  is  most  readily  obtained  by  beating  the  sulphate  with  carbon,  and  is  a 
whitish-yellow  mass.  It  melts  in  the  electric  furnace  and  crystallizes  on  cooling.  When 
it  is  dissolved  in  water  we  get  Calcium  Hydrosulphide,  Ca(SlI)2  (in  addition  to  cal- 
cium hydroxide),  which  decomposes  on  boiling  the  aqueous  solution.  When  calcium 
oxide  is  ignited  with  sulphur  in  a closed  crucible  a yellowish-gray  mass  is  obtained,  which 
consists  of  calcium  polysulphides  and  sulphate.  Milk  of  lime  boiled  with  sulphur  yields 
a deep-yellow  solution  of  calcium  polysulphides  and  calcium  hyposulphite.  When  these 
solutions  are  acted  upon  by  acids,  finely  divided  sulphur — milk  of  sulphur — is  precipi- 
tated and  hydrogen  sulphide  set  free.  If  the  reverse  occur,  viz.,  the  addition  of  a .solu- 
tion of  polysulphides  to  an  excess  of  dilute  acids,  hydrogen  persulphide  will  separate 
(p.  no). 

Calcium  Carbide,  CaCj.  Its  importance  in  illumination  has  been  dwelt  upon  on 
pp.  154,  155.  Fr.  Wohler  (1862)  first  described  it.  He  obtained  it  by  heating  intensely 
an  alloy  of  zinc  and  calcium  with  carbon  : 

ZnCa  -(-  Cj  — CaCg  -j-  Zn. 

W Borchers  was  the  first  person  to  prepare  it  in  the  electric  furnace.  In  1891  he  emphasized 
the  fact,  after  years  of  experience,  that  all  oxides  could  be  reduced  by  carbon  heated 
electrically.  Wilson  made  carbide  accidentally  on  a large  scale  in  1892.  Moissan  fol- 
lowed shortly  after.  Both  chemists  obtained  the  carbide  in  crystals.  It  maybe  procured 
pure  in  shining  gold-like,  non-transparent  crystals,  if  burnt  marble  is  heated  electrically 
with  sugar  carbon  : 

CaO  + 3C  = CaCj  + CO. 

Its  specific  gravity  is  2.2.  It  resists  most  of  the  strongly  reacting  substances,  yet  water 
and  dilute  acids  decompose  it  energetically  with  liberation  of  acetylene  (p.  154).  At 
present  it  is  made  in  many  factories  from  unslaked  lime  and  coal  or  coke  (see  the  publi- 
cation of  Liebetanz,  p.  155).  Strontium  and  barium  carbides  are  similar  to  calcium 
carbide. 


2.  STRONTIUM. 

Sr  = 87.6. 

This  element  is  rather  rare  in  nature,  and  is  principally  found  in  stron- 
tianite  (strontium  carbonate)  and  celestite  (strontium  sulphate).  Its 
compounds  are  very  similar  to  those  of  calcium. 

The  metal  is  obtained  by  the  electrolysis  of  fused  strontium  chloride. 
It  has  a brass-yellow  color,  and  a specific  gravity  of  2.5.  It  oxidizes  in 
the  air  and  burns  with  a bright  light  when  heated.  It  decomposes  water 
at  the  ordinary  temperature. 

Of  the  comj)ounds  of  strontium  we  may  mention  the  following: 

Strontium  Oxide,  SrO,  is  most  readily  obtained  by  igniting  the 
nitrate.  It  unites  with  water,  with  strong  evolution  of  heat,  forming 

Strontium  Hydroxide,  Sr(OH)2,  which  is  more  readily  soluble  in 
water  than  calcium  hydroxide.  It  crystallizes  from  aqueous  solution 
with  eight  or  nine  molecules  of  water.  When  ignited  it  decomposes  into 
strontium  oxide  and  water,  but  with  more  difficulty  than  calcium  hy- 
droxide. 

Strontium  Chloride,  SrCl2  -f-  6H2O,  crystallizes  from  water  in  hex- 
agonal tables,  which  deliquesce  in  the  air;  it  is  easily  soluble  in  alcohol. 

Strontium  Sulphate,  SrSO^,  is  much  more  difficultly  soluble  in  water 
than  calcium  sulphate,  but  is  not  as  insoluble  as  barium  sulphate  [one 
liter  dissolves  107  mg.  at 


3o8 


INORGANIC  CHEMIS^IRY. 


Strontium  Nitrate,  Sr(N()J^,  is  obtained  l)y  dissolving  tlic  carl)on- 
ate  in  nitric  acid,  and  is  rcatlily  soluble  in  water,  but  insoluble  in  alcohol. 
It  crystallizes  from  warm  solutions  in  anhydrous  octahedra,  but  from  cold, 
with  four  molecules  of  water,  in  monoclinic  i)risms.  Mixed  with  com- 
bustible substances  it  colors  the  flame  a beautiful  carmine-red,  and  for 
this  reason  is  employed  in  ji^  rotechny. 

Strontium  Carbonate,  SrCO.,,  is  [irecipitated  from  aqueous  solutions 
ol  strontium  salts  by  soluble  carbonates,  as  an  amorj)hous,  insoluble  pow- 
der. When  ignited  (iioo°)  it  breaks  down  into  strontium  oxide  and 
carbon  dioxide,  'bhis  decom])osition  does  not,  however,  occur  as  easily 
as  with  calcium  carbonate. 


8.  BARIUM. 

Ba  = 137.4. 

Barium  occurs  in  nature  in  large  masses,  as  heavy  spar  (or  barium  sul- 
phate),  and  as  witherite  (barium  carbonate).  All  its  compounds  are  dis- 
tinguished by  their  high  specific  gravity,  hence  the  name  barium,  from 
[■iafthq,  heavy.  In  accordance  with  its  general  character,  barium  is  a 
stronger  liasic  metal  than  either  strontium  or  calcium  (p.  299). 

The  barium  salts  are  either  prepared  from  the  natural  witherite,  by  dis- 
solving it  in  acids,  or  from  heavy  spar.  The  technical  procedure  con- 
sists in  heating  barium  sulphate  with  carbon  and  calcium  chloride  when 
barium  chloride  and  potassium  sulphide  (in  addition  to  carbon  monoxide) 
are  formed  : 

BaSO,  + 4C  + CaCb  = Bad,  + CaS  + 4CO. 

Metallic  barium,  like  strontium  and  calcium,  was  first  obtained  (Hum- 
phry Davy,  t8o8)  by  electrolysis.  Bunsen  irsed  the  fused  chloride  for  this 
purpose.  The  following  method  is  more  convenient : Sodium  amalgam 
is  added  to  a hot  saturated  barium  chloride  solution  ; the  sodium  dis- 
])laces  the  barium,  which  forms  an  amalgam  with  the  mercury.  The 
resulting  liquid  barium  amalgam  is  kneaded  with  water,  to  remove  all  the 
sodium,  and  then  heated  in  a stream  of  hydrogen,  to  volatilize  the 
mercury. 

Barium  is  a bright  yellow  metal,  of  specific  gravity  3.6.  It  fuses  at  a 
red  heat,  but  vaporizes  with  difficulty.  It  is  rapidly  oxidized  in  the  air; 
like  sodium  it  decomposes  water  very  energetically,  even  at  the  ordinary 
temjierature. 

Barium  Oxide,  BaO,  is  obtained  by  the  ignition  of  barium  nitrate. 
It  is  a gray,  amorphous  mass,  of  specific  gravity  5.5,  and  is  fusible  in  the 
oxyhydrogen  flame.  With  water  it  yields  the  hydroxide,  with  evolution 
of  much  heat. 

Barium  Hydroxide,  Ba(OH)2,  is  ])recipated  from  concentrated 
solutions  of  barium  salts  by  potassium  or  sodium  hydroxide,  not,  however, 
by  ammonium  hydroxide.  It  is  olitained  on  a large  scale  by  igniting 
heavy  spar  with  carbon,  when  barium  sulphide  is  formed,  and  its  aqueous 
solution  desulj)hurized  with  zinc  oxide: 

BaS  -|-  ZnO  | H,()  = ZiiS  + Ba(()H), ; 


BARIUM.  309 

also  Upon  heating  barium  carbonate  together  with  carbon  in  a current 
of  steam  : 

BaCOa  + C + lip  = Ba(OH)2  + 2CO. 

At  the  ordinary  temperature  it  dissolves  in  20  parts  of  water;  upon 
boiling,  in  3 parts.  From  aqueous  solution  it  crystallizes  with  eight  mole- 
cules of  water  in  four-sided  prisms  or  leaflets.  The  solution — called 
baryta  water — is  strongly  alkaline  and  is  very  similar  to  the  alkalies. 
When  exposed  to  the  air  it  absorbs  carbon  dioxide  and  becomes  turbid, 
with  separation  of  barium  carbonate.  At  a red  heat  it  fuses,  like  the 
caustic  alkalies,  without  decomposition,  and  on  cooling  solidifies  to  a 
crystalline  mass. 

Barium  Peroxide,  BaO.^,  is  produced  when  barium  oxide  (ignited 
barium  nitrate)  is  heated  in  a current  of  dry  air  (free  from  carbonic  acid) 
or  oxygen  ; it  always  contains  barium  oxide.  To  purify  it,  the  commercial 
peroxide  is  rubbed  together  with  water  and  added  to  cold,  very  dilute 
hydrochloric  acid,  until  the  latter  is  almost  saturated.  An  excess  of 
baryta  water  is  added  to  the  solution,  containing  barium  chloride  and 
hydrogen  peroxide,  when  hydrated  barium  peroxide,  BaO.^  + 8H2O,  sepa- 
rates in  shining  scales,  which,  upon  warming,  readily  lose  water  and 
break  down  into  a white  powder  consisting  of  barium  peroxide.  The 
latter  is  a compact  gray  mass  when  obtained  directly  from  the  oxide. 

The  peroxide  dissolves  in  dilute  acids,  with  production  of  hydrogen 
peroxide.  Concentrated  sulphuric  acid  sets  free  ozonized  oxygen  from 
it.  When  strongly  ignited  (at  about  700°)  it  decomposes  into  barium 
oxide  and  oxygen.  See  p.  81  for  its  use  in  the  technical  production  of 
oxygen.  It  is  also  employed  for  bleaching  purposes  and  for  the  prepara- 
tion of  hydrogen  peroxide;  see  pp.  99,  100. 

Barium  Chloride,  BaCl2,  crystallizes  from  aqueous  solution,  with 
two  molecules  of  water,  in  large,  rhombic  tables,  which  are  stable  in  the 
air.  It  dissolves  readily  in  water,  and  is  poisonous,  like  all  soluble 
barium  salts. 

Barium  Nitrate,  Ba(N03)2,  crystallizes  in  anhydrous,  shining  octa- 
hedra,  of  the  regular  system,  soluble  in  12  parts  of  cold  and  3 parts  of 
hot  water.  Technically  it  is  made  by  mixing  concentrated  solutions  of 
barium  chloride  with  sodium  nitrate,  when  barium  nitrate  is  precipitated. 
It  is  employed  as  a green  fire  in  pyrotechny. 

Barium  Sulphate,  BaSO^,  is  found  in  nature  as  heavy  spar,  in 
rhombic  j)risms,  with  a specific  gravity  of  4.6.  It  is  obtained  artificially 
by  the  precipitation  of  barium  salts  with  sulphuric  acid,  in  the  form  of  a 
white,  amorphous  powder,  almost  insoluble  in  water  and  acids.  Under 
the  name  of  permanent  white,  it  is  used  as  a paint,  as  a substitute  for  poi- 
sonous white  lead,  from  which  it  is  also  distinguished  by  its  unaltera- 
bility. 

Barium  Persulphate,  BaS.^O^  fpp.  188,  198),  obtained  by  rubbing  the  ammonium 
salt  together  with  barium  hydroxide,  crystallizes  from  its  solution  {in  vacuo)  in  prisms, 
containing  four  molecules  of  water  of  crystallization.  It  soon  decomposes,  yielding 
barium  sulphate,  suh)huric  acid,  and  oxygen  : 

BaSPg  -f  II2O  = BaSO^  i 11,80^  -p  O. 


310 


INORGANIC  CHEMISTRY. 


Barium  Carbonate,  IkiCO.,,  as  witherite,  occurs  iu  sliiuing,  rlioml)ic 
crystals,  and  is  i)recipilatc(l  from  barium  solutions  by  soluble  caibonatcs, 
as  a white,  amorphous  i)()wder.  It  fuses  at  a white  heat,  and  loses  carbon 
dioxide  at  1400-1500°. 

Barium  Sulphide,  llaS,  is  obtained  l)y  igniting  the  sul|>hate  with 
carbon.  It  dissolves  in  water,  with  decomjiosition  into  hydroxide  and 
hydrosulphide. 


RECOGNITION  OF  THE  COMPOUNDS  OF  THE  ALKALINE 

EARTHS. 

The  carbonates  and  phosphates  of  this  group  are  insoluble  in  water; 
hence  are  precipitated  from  the  aqueous  solutions  of  their  salts  upon  the 
addition  of  soluble  carbonates  and  phosphates  (of  the  alkalies).  The 
sulphates  are  also  sparingly  soluble  in  acids;  for  this  reason  they  are 
thrown  down  from  acid  solutions  l^y  soluble  sulphates  or  free  sulf)huric 
acid;  the  precipitation  is  complete,  even  with  calcium,  if  alcohol  be 
added  to  the  solution,  d'he  hydroxides  of  the  alkaline  earths,  which  are 
more  or  less  soluble  in  water,  are  only  precipitated  by  sodium  or  potas- 
sium hydroxide  from  concentrated  solutions.  In  solutions  of  barium 
salts  hydrofluosilicic  acid  produces  a crystalline  y^reciju'tate  of  barium 
silicofluoride,  BaSiFlg,  and  potassium  chromate  one  of  barium  chromate, 
BaCrO,. 

Koblrausch  and  Rose  determined,  by  means  of  the  electric  conductivity  of  solutions, 
the  solubility  of  many  of  the  salts  which  have  been  mentioned — the  so-called  insoluble 
salts.  They  found  that  a liter  of  water  at  15°  dissolved  o.  i mg.  of  silver  iodide  ; 0.4  mg. 
of  silver  bromide  ; 0.5  mg.  of  mercuric  iodide  ; 1.7  mg.  of  silver  chloride  ; 3.1  mg.  of  mer- 
curous chloride  ; 2.6  mg.  of  barium  sulphate  ; 107  mg.  of  strontium  sulphate  ; 2070  mg. 
of  calcium  sulphate  ; 13  mg.  of  calcium  carbonate  ; 24  mg.  of  barium  carbonate  ; 3.8  mg. 
of  barfum  chromate;  0.2  mg.  of  lead  chromate;  9 mg.  of  magnesium  hydroxide.  See 
Z.  f.  anorg.  Ch.  5 (1894),  237. 

The  flame  colorations  produced  by  the  volatile  compounds  are  very 
characteristic;  calcium  salts  impart  a reddish-yellow  color;  strontium, 
an  intense  crimson  ; barium,  a yellowish-green.  The  spectrum  of  calcium 
exhibits  several  yellow  and  orange  lines,  and  in  addition,  a green  and 
a violet  line  (see  the  Spectrum  Table);  that  of  strontium  contains, 
besides  several  red  lines,  an  orange  and  a blue,  which  are  less  distinct 
but  very  characteristic.  Finally,  the  barium  spectrum  consists  of  several 
orange,  yellow,  and  green  lines,  of  which  a bright  green  is  particular!  \ 
prominent. 


DIAMMONIUM  COMPOUNDS. 

Tlie  .same  relation  which  ammonia  sustains  to  the  univalent  alkali  metals  (p.  296), 
hvtirazine  or  dianiide,  N^II^  (pp.  125,  131  ),  discovered  by  Curtins,  bears  to  the  bivalent 
alkaline  earth  metals.  Diamide,  however,  does  not  always  act  as  a di-acid  base,  which 
might  be  inferred  from  what  has  been  said,  for  in  its  most  stable  derivatives  it  is  as  mon- 
acid  as  ammonia.  It  may  be  assumed  that  the  di-acid  compounds  contain  the  bivalent 
II 

radical  diainnioniiiin,  1^2*  counterpart  of  the  radical  ammonium,  and  that  they 


DIAMMONIUM  COMPOUNDS. 


3II 


accordingly  arrange  themselves  with  the  salts  of  barium,  strontium  and  calcium.  The 
similarity  of  diammonium  to  the  alkaline  earth  metals  is  shown  also  in  the  difficult  solu- 
bility of  the  sulphate  and  by  its  inability  to  form  alums. 

Again,  diamide  manifests  properties  allying  it  with  the  alkali  metals.  Its  hydrate, 
N.^H^ . 11.^0  = N.^IIgOIl,  generally  behaves  like  a mon-acid  base.  Its  chloride,  N2HgCl,^, 
breaks  down  below  100°  into  hydrochloric  acid  and  the  chloride  N.^HgCl,  which  does  not 
yield  any  more  hydrochloric  acid  unless  destroyed  by  heating.  The  hydrate  N2H4  . - 
2H2O  = N2Hg(OH)2  is  only  stable  in  aqueous  solution  and  upon  evaporation  changes  to 
N2ll4.  H2O  = N2II5OM,  boiling  without  decomposition.  It  forms  but  one  nitrate  of  the 
composition  N2H4.HN0.4,  but  one  sulphocyanide  N2II4.HCNS,  and  but  one  diam- 
monium nitride,  N2H4 . N3H. 

It  would  therefore  appear  that  both  the  bivalent  radical  N2Hg  and  the  univalent  radical 
NjHg  are  present  in  the  compounds  of  diamide,  and  further  that  those  containing  the 
radical  N2H5  are  the  more  stable  derivatives. 


f OH 


Dihydrate. 


N2H5OH 

Mono-hydrate. 


Unstable  Diammonium  Compounds. 

Dichloride.  Di-iodide. 

Stable  Diammonium  Compounds. 

N2H5CI  N2H5I 

Mono-chloride.  Mono-iodide. 

N.HgNOg. 

Nitrate. 


N2HgS04. 

Sulphate. 

(N2Hg)2S04 

Semi-sulphate. 


The  hydrazine  .salts  generally  crystallize  well.  Their  great  reducing  power  is  one 
of  their  characteristics.  They  precipitate  metallic  silver  from  ammoniacal  silver  solutions  ; 
cuprous  oxide  and  metallic  copper  from  alkaline  copper  solutions  (Fehling’s  solution),  and 
gold  from  acid  solutions  of  gold  chloride  (distinction  from  hydroxylamine).  The  hydra- 
zine salts  are  very  poisonous.  With  care  nitrous  acid  will  liberate  hydrazoic  acid  from 
them  : 

N2H4  + HNO2  = N3H  -f  2H2O, 


otherwise  nitrogen.  The  addition  of  copper  sulphate  to  the  solutions  of  hydrazine  salts 
causes  the  separation  of  a sparingly  soluble  double  salt  CUSO4  . (N2H-).2S04  ; this  can  be 
used  to  separate  hydrazine  from  its  solutions  and  from  mixtures.  The  corresponding 
double  salt.s,  with  the  salts  of  nickel,  cobalt,  iron,  manganese,  cadmium  and  zinc,  are  also 
soluble  with  difficulty  and  are  anhydrous.  P'or  the  detection  by  means  of  benzaldehyde 
(benzalazine,  m.  p.  93°)  see  Organic  Chemistry.  Consult  Curtius  and  Schrader,  Jr. 
prakt.  Ch.  50  (1894),  31 1. 

II 

Diammonium  Sulphate,  N2II4.II2SO4  = (N2Hg)S04,  can  be  made  by  treating 
solutions  of  all  the  other  hydrazine  salts  with  sulphuric  acid.  It  consist  of  vitreous  plates, 
which  are  soluble  in  about  33  parts  of  water  at  20°.  They  deflagrate  when  quickly  heated. 

II 

Diammonium  Chloride,  N2II4.  2IICI  = (N2Hg)Cl2,  obtained  by  transposing  the 
sulphate  with  barium  chloride,  forms  large,  vitreous  octahedra.  It  is  very  soluble  in 
water.  It  melts  with  decomposition  at  about  200°.  At  100°  it  passes  gradually  into 
monohydrate. 

I 

Hydrazine  Hydrate,  (diammonium  monohydrate),  N2II4.  = (N2ll5)OII,  can 
be  obtained  by  distilling  hydrazine  sulphate  with  caustic  potash.  It  is  a strongly  refract- 
ing, not  very  mobile  liquid,  which  fumes  in  the  air  and  boils  at  118.5°  under  739.5  mm. 
pres.sure.  It  has  a faint  odor,  not  reminding  one  of  ammonia,  has  a caustic  taste,  is 
hygro-scojjic  and  attracts  carbon  dioxide  from  the  air.  Below  — 40°  it  solidifies  to  a 
leafy  crystalline  mass.  Its  specific  gravity  equals  1.03  at  21°.  It  sinks  in  water  and 
only  mixes  with  it  after  some  time.  It  can  be  preserved  in  a pure  condition  and  in  con- 
centrated solution.  Dilute  aqueous  .solutions  decompose  completely  in  (he  course  of 
time.  Its  .solution  and  its  vapors  color  red  litmus  blue.  It  breaks  dowm  (from  its  vapor 
density)  at  171°  and  760  mm.  jire.ssure  into  hydrazine  and  water,  but  on  the  other  hand 
it  vaporizes  in  a vacuum  at  100°  without  decomposition.  It  has  a powerful  reducing 


312 


INORGANIC  CHKMISTKY. 


action  upon  many  oxides  ; with  cliromic  acid  and  mercuric  oxide  it  causes  explosion. 
Nothing  delinite  is  known  as  to  its  chemical  constitution.  [See  Curtins  and  .Schul/,  Jr. 
j)rakt.  Ch.  42  (1890),  521  ; Curtins  and  Schrader,  ibid.  50  (1894),  318J.  Upon  distill- 
ing the  hydrate  with  barium  oxide  under  reduced  pressure  the  diamide  boiling  only  at 
lower,  is  set  free.  It  fumes  strongly  in  the  air  and  is  combustible. 

I 

Diammonium  Monochloride,  NjH^.  HCl  _ (N.^Il5)Cl,  results  when  the  dichloride 
is  heated.  It  crystallizes  in  long  white  needles,  melting  at  89°.  It  is  very  soluble  in 
water. 

I 

Diammonium  Semisulphate,  ( N^I  I^).^ . I I.^SD^  — (N^n.),^SO,,  is  formed  ui)on 
neutralizing  hydrazine  hydrate  with  sulphuric  acid.  It  crystallizes  in  large,  delicjuesccnt 
plates,  melting  at  85°. 

Another  reTnarkable  body  is  the  Diammonium  Nitride,  N5II5  — N,^lI^.N3ll  = 

(N2ll5)N.j,  prepared  by  Curtins  from  ammonium  nitride  and  hydrazine  hydrate.  It  is  the 

II 

ammonium  salt  of  hydrazoic  acid.  'I'he  other  salt,  Ngl  1^,  = K^If^ . fN^II).^  = (N^I  I,.)(N.j)2, 
corresponding  to  the  dichloride  described  above,  could  not  be  obtained.  Diaminonium 
nitride  (diaminonium  monazide,  if  the  salts  of  hydrazoic  acid  are  designated  azides)  con- 
sists of  large,  vitreous  prisms,  melting  at  about  50°,  deliijuescing  in  the  air  and  volatiliz- 
ing at  the  ordinary  temperature,  but  more  readily  with  acpieous  and  alcoholic  vapors, 
d'he  salt  dissolves  with  difficulty  in  alcohol.  When  ignited  it  burns  cpiietly  with  a yellow 
flame,  without  smoke  and  without  residue.  If  heated  rapidly  in  the  air,  by  contact  with 
a wire  raised  to  a white  heat  by  ignition  with  detonating  metallic  azides  or  fulminating 
mercury,  a powerful  explosion  will  take  place  even  if  the  salt  has  deliquesced  [Ber.  24 
(1891),  3348  ; see  also  Ammonium  Nitride,  p.  298.] 


2.  METALS  OF  THE  MAGNESIUM  GROUP. 

Id'  this  group  are  usually  included  beryllium,  magnesium,  zinc,  and 
cadmium.  However,  these  metals  do  not  exhibit  complete  analogy,  as  is 
clearly  seen  from  their  position  in  the  periodic  system  (p.  300).  Beryl- 
lium shows  the  greatest  variations.  It  approaches  aluminium,  while  mag- 
nesium resembles  not  only  zinc  and  cadmium,  but  also  the  alkaline  earth 
metals,  calcium,  strontium,  and  barium.  Its  similarity  to  the  latter  is 
expressed  by  the  basic  nature  of  its  oxide,  whereas  it  resembles  zinc  and 
cadmium  mainly  in  isomorphism  of  compounds. 

Beryllium  and  magnesium  bear  the  same  relations  to  calcium,  stron- 
tium, and  barium,  as  lithium  and  sodium  bear  to  the  metals  of  the  potas- 
sium grou]:). 

The  alkaline  character  of  the  alkaline  earth  metals  gradually  diminishes 
from  barium  to  calcium,  and  becomes  almost  nothing  in  magnesium  and 
beryllium,  which  possess  the  lowest  atomic  weights  (see  p.  300).  Mag- 
nesium and  beryllium  are  scarcely  capable  of  decomposing  water,  even  at 
l)oiling  temperatures.  Their  oxides  and  hydroxides  are  almost  insoluble 
in  it  ; the  hydroxides  decompose,  on  gentle  ignition,  into  oxides  and 
water.  'I'lieir  carbonates  are  very  unstable;  their  chlorides,  too,  suffer  a 
partial  decom])osition  into  oxide  and  liN’drogen  chloride,  even  on  drying, 
d'lie  solubility  of  the  sulphates  of  magnesium  and  beryllium  further  dis- 
tinguishes them  from  the  metals  of  the  alkaline  earth  group. 


MAGNESIUM. 


3^3 


The  specific  properties  of  beryllium  and  magnesium  are  maintained  in 
zinc  and  cadmium,  which  constitute  a natural  group  with  the  former. 
Zinc  and  cadmium  do  not  decompose  water  at  a boiling  heat;  their 
hydroxides  are  insoluble  in  it,  and  are  not  very  stable;  their  carbonates 
and  chlorides  easily  undergo  decomposition;  their  sulphates  are  readily 
soluble  in  water.  The  similarity  is  further  expressed  by  the  isomorphism 
of  most  of  their  compounds.  'I'hus,  magnesium  and  zinc  sulphates'crys- 
tallize  with  seven  molecules  of  water,  in  perfectly  similar  forms.  If  the 
solution  of  a mixture  of  both  salts  be  allowed  to  crystallize,  we  get  crystals 
with  variable  quantities  of  zinc  and  magnesium : the  formation  of  such 
tsomorphous  mixtures  in  ad  libitu7n  proportions  is  a characteristic  indica- 
tion of  the  isomorphism  of  compounds  chemically  similar  (p.  255). 

The  difference  between  beryllium  and  magnesium  upon  the  one  side, 
and  zinc  and  cadmium  on  the  other,  is  shown  distinctly  in  their  specific 
gravities.  While  the  first  two  elements  possess  a low  specific  gravity 
(Be  1.8,  Mg  1.74),  zinc  and  cadmium  (with  specific  gravities  of  7.1 
and  8.6)  belong  to  the  so-called  heavy  metals  (see  p.  252). 

The  difference  in  specific  gravity  determines,  also,  many  differences  in 
chemical  character.  The  light  metals  (especially  the  alkalies  and  alkaline 
earths)  form  rather  unstable  sulphides,  readily  soluble  in  water,  while  the 
sulphides  of  zinc  and  cadmium,  like  those  of  all  heavy  metals,  are  insolu- 
ble in  water,  and  usually  in  dilute  acids;  hence  magnesium  and  beryl- 
lium are  not,  while  zinc  and  cadmium  are,  precipitated  by  hydrogen  sul- 
phide or  alkaline  sulphides  as  sulphides  from  solutions  of  their  salts. 
Further,  the  oxides  of  the  light  metals  are  very  stable,  and  are  only 
reduced  by  ignition  with  carbon  if  they  are  readily  fusible  (like  potassium 
and  sodium  oxides);  the  heavy  metals,  on  the  other  hand,  are  easily  separ- 
ated from  their  oxides  by  carbon.  Zinc  and  cadmium  oxides  are  reduced 
by  carbon,  while  those  of  magnesium  and  beryllium  are  not  altered. 


1.  MAGNESIUM. 

Mg  = 24.36. 

Magnesium  is  very  abundant  in  nature,  and  almost  always  accompanies 
calcium  in  its  compounds.  As  carbonate,  it  occurs  in  compact  masses, 
as  magnesite,  etc.  Dolomite,  which  forms  entire  mountains,  is  an  iso- 
morphous  mixture  of  calcium  and  magnesium  carbonates.  Magnesium 
is  also  present  in  most  of  the  natural  silicates;  its  soluble  salts  are  con- 
tained in  almost  all  natural  waters.  In  conjunction  with  the  alkali  salts 
they  constitute  the  most  important  salts  of  the  Stassfurt  beds  (see  p. 
276). 

Metallic  magnesium  may  be  obtained  by  the  electrolysis  of  the  chloride 
(Bunsen,  1852)  or  by  heating  the  same  with  sodium  (Bussy,  1830).  It  is 
more  easily  prepared  by  heating  the  double  chloride  of  magnesium  and 
sodium  with  metallic  sodium  and  fluorspar,  the  latter  serving  merely  as  a 
flux  : 

MgCb.XaCl  -f-  2Na  = 3NaCl  4 Mg. 

27 


3^4 


INORGANIC  CHKMlh'I'RV. 


At  present  magnesium  is  ol)tained  in  large  (luantities  by  the  electrol- 
ysis of  fused  carnallite  [see  Dammer,  Chemische  'Fechnologie,  Ikl.  ii, 

(1895)]- 

Magnesium  is  a brilliant,  almost  silver-white  metal,  of  specific  gravity 
1.74.  It  is  tenacious  and  ductile,  and  when  heated  may  be  converted 
into  wire  and  rolled  out  into  thin  ribbons.  It  melts  at  about  700°,  and 
distils  at  a bright-red  heat.  At  ordinary  temperatures,  it  scarcely  oxidizes 
in  the  air;  it  burns,  wlien  heated,  with  an  extremely  intense  white  light, 
owing  to  the  glowing  non-volatile  magnesium  oxide.  Magnesium  light 
is  rich  in  chemically  active  rays,  and,  for  this  reason,  it  is  employed  in 
photograi)hy  for  artificial  illumination.  Its  alloy  with  zinc  is  generally 
em})loycd  as  a substitute  for  pure  magnesium,  as  it  burns  with  an 
equally  bright  light.  Its  intense  light  has  led  to  its  use  in  pyrotechny. 
Boiling  water  is  very  slowly  decom])osed  by  magnesium.  It  dissolves 
easily  in  dilute  acids,  forming  salts;  the  alkalies  do  not  attack  it.  Many 
metallic  oxides  and  acids  when  heated  with  magnesium  powder  lose  their 
oxygen  and  are  reduced  by  it  (see  p.  237). 

Magnesium  Oxide,  MgO,  or  ;/n7^i^;/esia,  formed  by  the  combustion 
of  magnesium,  is  ordinarily  obtained  by  the  ignition  of  the  hydroxide  or 
the  carbonate  (^Magnesia  iista).  It  is  a white,  very  voluminous,  amor- 
phous powder,  which  finds  apjdication  in  medicine,  d'he  feebly  ignited 
magnesia  combines  with  water,  with  slight  generation  of  heat,  to  produce 
magnesium  hydroxide. 

In  Stassfurt  the  magnesium  chloride  liciuors,  produced  in  the  prepara- 
tion of  potassium  chloride,  are  utilized  in  j^roducing  a very  good  mag- 
nesia, which  is  used  in  the  manufacture  of  fire-brick  or  fire-clay  (p.  276). 

At  high  temperatures  magnesia  conducts  electricity  and  it  can  be  made  to  glow  and 
become  luminous  by  the  current.  If,  therefore,  little  thin  cylinders  of  magnesia  are 
heated  until  they  reach  the  point  of  conductivity  they  will  answer  for  the  production  of 
the  electric  light  (Nernst). 

Magnesium  Hydroxide,  Mg(OH)2,  is  precipitated  from  solutions 
of  magnesium  salts  by  potassium  or  sodium  hydroxide  as  a gelatinous 
mass.  Dried  at  100°  it  is  a white  amorphous  powder.  It  is  almost  in- 
soluble in  water  and  alkalies;  moist  red  litmus-paper  is,  however,  colored 
blue.  Ammonium  salts  dissolve  it  quite  easily,  forming  soluble  double 
salts.  Magnesium  hydroxide  attracts  carbon  dioxide  from  the  air  and 
forms  magnesium  carbonate.  It  yields  the  oxide  and  water  when  gently 
ignited. 

Magnesium  Chloride,  MgCl2,  is  present  in  traces  in  many  mineral 
s])rings.  It  may  be  prepared  by  the  solution  of  the  carbonate  or  oxide 
in  hydrochloric  acid  ; in  large  quantities  it  is  obtained  as  a by-])roduct 
in  the  technical  ])roduction  of  potassium  chloride.  When  its  solution 
is  evaporated  the  salt  crystallizes  out  with  six  molecules  of  water  in  deli- 
quescent crystals,  isomori)hous  with  calcium  chloride.  When  these  are 
heated  they  give  up  water,  and  there  occurs  at  the  same  time  a partial 
decomj)osition  of  the  chloride  into  oxide  and  hydrogen  chloride: 


MgC:i2  1 11,0  MgO  4-  2IIC1. 


MAGNESIUM  SULPHATE. 


315 


As  magnesium  chloride  is  produced  in  large  quantities  in  various  technical  processes, 
repeated  efforts  liave  been  made  to  utilize  the  above  reaction  for  the  preparation  of  hydro- 
chloric acid,  by  conducting  steam  over  heated  magnesium  chloride. 

Lately  the  liquors  have  been  evaporated  and  the  magnesium  chloride  dehydrated, 
whereby  it  breaks  down  in  part  into  hydrochloric  acid  and  magnesium  oxide.  When  the 
residual  magnesium  oxychloride  is  heated  in  an  air  current,  magnesia  and  chlorine  are 
produced  (Weldon-Pechiney  ; see  p.  50). 

To  get  anhydrous  magnesium  chloride,  ammonium  chloride  is  added  to 
the  solution  of  the  former.  The  double  salt,  MgCl2.NH^Cl  -|-  6H2O, 
is  formed.  When  this  is  heated  it  first  loses  water,  and  at  460°  throws 
off  ammonium  chloride,  leaving  anhydrous  magnesium  chloride.  The 
latter  can  also  be  obtained  by  heating  the  hydrated  salt  in  a current  of 
hydrochloric  acid.  This  is  a leafy,  crystalline  mass,  which  fuses  at  708°, 
and  distils  undecomposed  at  a red  heat;  it  is  very  deliquescent  in  the  air. 

Double  salts,  similar  to  the  above,  are  also  formed  from  potassium  and 
calcium  chloride.  The  potassium  double  salt,  MgCl2.KCl  -f-  6H2O, 
occurs  in  considerable  deposits  as  carnallite  at  Stassfurt  (p.  276). 

Magnesium  Sulphate,  MgSO^,  is  found  in  sea-water  and  in  many 
mineral  springs.  With  more  or  less  water  it  is  kieserite^  which  abounds 
extensively  at  Stassfurt.  At  ordinary  temperatures  it  crystallizes  with 
seven  molecules  of  water,  MgSO^  -j-  7H2O,  in  four-sided  rhombic  prisms, 
readily  soluble  in  water  (at  0°  in  2 parts  of  w’ater).  It  has  a bitter,  salt-like 
taste,  and  serves  as  an  aperient.  It  crystallizes  with  six  molecules  of  water 
'from  solutions  heated  to  70°;  at  0°,  however,  it  has  twelve  molecules. 
When  heated  to  150°  these  hydrates  lose  all  their  water  of  crystallization, 
excepting  one  molecule,  which  escapes  above  200°.  One  molecule  of 
water,  in  magnesium  sulphate,  is,  therefore,  more  closely  combined  than 
the  rest.  Many  other  salts  containing  water  deport  themselves  similarly. 
The  more  intimately  combined  water  is  termed  water  of  constitution. 

Magnesium  sulphate  forms  double  salts  with  potassium  and  ammonium 
sulphates,  which  crystallize  with  six  molecules  of  water  in  monoclinic 
prisms,  e.  g : 

MgSO,.K2SO,  4-  6H2O. 

The  sulphates  of  zinc  and  several  other  metals,  e.  g.,  iron,  cobalt,  and  nickel,  in  their 
bivalent  forms,  are  very  similar  to  magnesium  sulphate.  Their  sulphates  crystallize  with 
seven  molecules  of  water,  and  are  isomorphous.  They  form  double  salts  with  potassium 
and  ammonium  sulphates  ; these  crystallize  with  six  molecules  of  water,  and  are  iso- 
morphous, e.  g.  : 

ZnSO,  + 7H2O  ZnSO, . + 6H2O. 

FeSO^  -b  7H2O  FeSO^.  K2SCh  -f  6Uf. 

The  constitution  of  these  double  salts  may  be  viewed  in  the  same  way  as  that  of  potas- 
sium-sodium sulphate,  or  of  mixed  salts  of  polybasic  acids.  We  may  suppose  that  in  the 
given  instance  the  bivalent  metal  unites  two  molecules  of  sulphuric  acid  : 

II 

^04/Mg  -f6H20. 

Magnesium  Phosphates. — The  tertiary  phosphate  Mg,(POJ2> 
accompanies  the  tertiary  calcium  phosphate  in  small  quantities  in  bones 


INORGANIC  CHEMISTRY. 


316 

and  in  plant  ashes.  The  secondary  phosphate,  MgHPO^  -|-  is 

precipitated  from  the  soluble  magnesium  salts,  by  disodium  phosphate 
(NajHPO^)  as  a salt  dissolving  with  difficulty  in  water,  in  the  pres- 
ence of  ammonia  and  ammonium  salts,  magnesium-ammonium  phosphate , 
MgNH^PO^  -f-  ^H/),  is  precipitated  as  a crystalline  powder  insoluble  in 
water.  This  double  salt  is  found  in  guano,  forms  in  the  decay  of  urine, 
and  is  sometimes  the  cause  of  the  formation  of  calculi.  The  primary 
salt,  MgHj(PO^)2,  has  not  been  obtained. 

The  magnesium  salts  of  arsenic  acid,  H^AsO^,  are  very  similar  to  those 
of  phosphoric  acid.  Magnesium-ammonium  arseniate,  MgNH^AsO^ -f- 
6H2O,  is  likewise  almost  insoluble  in  water. 

Magnesium  Carbonate,  MgCO,,  occurs  in  nature  as  magnesite, 
crystallized  in  rhombohedra  ( isomorphous  with  calcite),  or  more  often 
in  compact  masses.  Combined  with  calcium  carbonate,  it  forms  dolo- 
mite, to  which,  when  pure,  is  ascribed  the  formula,  CaCO^.MgCOg; 
however,  it  usually  contains  an  excess  of  calcium  carbonate.  On  adding 
sodium  or  potassium  carbonate  to  the  acpieous  solution  of  a magnesium 
salt,  some  carbon  dioxide  escapes,  and  a white  precipitate  forms,  which 
consists  of  a mixture  of  magnesium  carbonate  and  hydroxide.  If  the 
precipitate  be  dried  at  low  temperature,  we  obtain  a white,  voluminous 
powder,  whose  composition  generally  corresponds  to  the  formula 
Mg(OH)2.  3MgC03 -f- 4H,^0.  'Phis  salt  is  em[)loyed  under  the  name 
Magnesia  alba  in  medicine.  If  it  be  suspended  in  water,  and  carbon 
dioxide  passed  through  it,  the  salt  will  dissolve,  and  upon  standing 
exposed  to  the  air,  crystals  of  neutral  cai'bonate,  MgCOg  3^2^?  sepa- 
rate. When  these  are  boiled  with  water  they  give  up  carbon  dioxide 
and  are  again  converted  into  the  basic  carbonate.  The  naturally  occur- 
ring magnesite  sustains  no  change  when  boiled  with  water,  and  it  is 
only  when  it  is  heated  above  300°  that  it  decomposes  into  magnesium 
oxide  and  carbon  dioxide. 

Freshly  prepared  magnesium  carbonate  dissolves  in  acid  alkaline  car- 
bonates and  crystallizes  from  them  as  4MgC03  -{-  15H2O. 

Magnesium  carbonate  yields  isomorphous  double  salts,  with  potassium 
carbonate  and  ammonium  carbonate  ; e.  g.,  MgC03.K2C03  -f-  4H2O. 

Of  the  silicates  of  magnesium,  we  may  mention  olivine  (Mg2SiO^), 
serpentine  (Mg3Si207  -f"  2H2O),  talc  (Mg^Si^O^j  -f-  H2O),  sepiolite  or 
meerschaum  (Mg2Si30g  -f-  2H2O).  The  mixed  silicates  of  magnesium 
and  calcium  are  very  numerous;  to  these  belong  asbestos,  the  augites  and 
horn-blendes. 


Magnesium  Nitride,  Mg.^Nj  (pp.  116,  124,  126),  is  a light,  porous,  yellow-colored 
mass.  liriegleb  and  (ieuther  found  that  it  was  produced  when  magnesium  was  heated 
in  a current  of  ammonia  : 

Mg3  t 2NIl3  = Mg3N, -f  3lb. 

It  is  also  obtained  pure  when  magnesium  is  heated  in  nitrogen.  When  introduced  into 
water,  the  latter  is  made  to  boil  and  large  volumes  of  ammonia  escape  [see  Pasch- 
kowezky,  Jr.  prakt.  Ch,  47  (1893),  89]. 


BERYLLIUM. 


317 

Recognition  of  Magnesium  Compounds. — The  fixed  alkaline  hy- 
droxides precipitate  magnesium  hydroxide  from  soluble  magnesium  salts; 
the  carbonates  throw  down  basic  magnesium  carbonate.  The  precipitates 
are  insoluble  in  pure  water  and  the  alkalies,  but  dissolve  readily  in  solu- 
tions of  ammonium  salts.  In  the  presence  of  the  latter,  neither  the  alka- 
line hydroxides  nor  carbonates  cause  precipitation.  In  the  presence  of 
ammonia  and  ammonium  chloride,  disodium  phosphate  precipitates  mag- 
nesium-ammonium phosphate,  MgNH^PO^  insoluble  in  water. 


2.  BERYLLIUM. 

Be  = 9.1. 

Among  the  metals  of  the  second  group  beryllium  occupies  a position  similar  to  that  of 
lithium  in  the  first  group  (pp.  272,  299)  ; in  both  elements,  which  have  the  lowest  atomic 
weight  in  their  group,  the  specific  group  character  is  considerably  diminished,  or  does  not 
find  expression.  As  lithium  attaches  itself  in  many  respects  to  magnesium,  so  does  beryl- 
lium approach  aluminium.  Like  the  latter,  it  is  scarcely  at  all  attacked  by  nitric  acid,  but 
dissolves  easily  in  sodium  or  potassium  hydroxide,  with  elimination  of  hydrogen.  Like 
aluminium  oxide,  that  of  beryllium  dissolves  in  the  alkalies,  and  is  almost  invariably 
accompanied  by  the  former  in  its  natural  compounds.  However,  beryllium,  in  most  of 
its  compounds,  stands  nearer  to  magnesium  than  to  aluminium.  The  determination  of 
the  vapor  density  of  beryllium  chloride  (see  below)  and  of  certain  organic  derivatives  has 
finally  established  the  atomic  weight  and  the  valence  of  this  element. 


Beryllium  is  not  very  abundant  in  nature  and  is  found  principally  in  beryl,  a double 
silicate  of  aluminium  and  beryllium,  Al2Be3(Si03)(,.  Emerald  has  the  same  composi- 
tion, and  is  only  colored  green  by  a slight  amount  of  chromium  oxide. 

It  was  in  these  minerals  that  V auquelin  discovered  beryllium  in  1797.  It  is  also  present 
in  leucophane  (a  silicate),  together  with  aluminium,  fluorine  and  sodium  ; and  in  gado- 
linite  (a  silicate)  with  ferrous  oxide,  yttrium,  cerium,  lanthanum,  and  other  rare  earths. 

Metallic  beryllium  is  obtained  by  the  ignition  of  the  chloride,  or  better,  potassium  beryl- 
lium chloride  (or  potassium  beryllium  fluoride,  BeFl2.2KFl),  with  sodium,  and  by  the 
electrolysis  of  BeFlg.  KFl.  It  is  a white,  ductile  metal,  of  specific  gravity  1.8.  Its  specific 
heat  at  the  ordinary  temperature  equals  0.408  ; the  atomic  heat  is,  therefore,  3.7 
(p.  254).  It  does  not  decompose  water,  even  upon  boiling.  It  does  not  oxidize  in  the 
air  at  ordinary  temperatures.  When  finely  divided  and  heated  it  will  burn  in  the  air 
with  a very  bright  light.  It  is  readily  dissolved  by  dilute  hydrochloric  and  sulphuric 
acids  ; also  by  potassium  and  sodium  hydroxides. 

Beryllium  Chloride,  BeCl2,  is  obtained,  like  aluminium  chloride,  by  the  ignition  of 
a mixture  of  beryllium  oxide  and  carbon  in  a stream  of  chlorine.  It  sublimes  in  shin- 
ing needles,  which  deliquesce  in  the  air.  Its  vapor  density  corresponds  to  the  molecular 
formula  BeCl2  ==  79.9  (Nilson).  Pure  beryllium  chloride  dissolves  in  water,  forming  a 
colorless  solution,  from  which  it  crystallizes  with  four  molecules  of  water;  upon  drying  it 
suffers  a decomposition  similar  to  that  of  magnesium  chloride. 

The  salts  of  beryllium  have  a sweet  taste,  hence  it  has  been  called  glticinwn. 
Ammonium  hydroxide  precipitates  a white,  gelatinous  beryllium  hydroxide,  Be(OH)2, 
from  solutions  of  the  soluble  salts.  This  dissolves  readily  in  sodium  and  potassium 
hydroxide,  and  in  a mixture  of  ammonia  and  ammonium  carbonate,  but  on  boiling,  sepa- 
rates again  from  solution.  When  heated,  the  hydroxide  breaks  down  into  water  and 
beryllium  oxide,  Bef),  which  is  a white,  amorjflious  powder,  of  specific  gravity  2.96. 

Beryllium  Sulphate,  BeSO^,  crystallizes  from  water  at  various  temperatures,  with 
four  or  seven  molecules  of  water ; it  contains  two  molecules  at  105°,  and  is  anhydrous  at 


INORGANIC  CIIKMIS'I'RY. 


318 

250-260°,  It  docs  not  crystallize  with  inagiicsiuin  siilpliatc  in  an  isoinorphous  niixlnrc. 
d'he  double  salt,  HeSO^.  K.^St  >4  -)-  2ll2<),  docs  not  dissolve  readily  in  water.  Compare 
Kruss  and  Moraht,  Her.  23  (1890),  727,  and  Ann.  Chein.  262  (1891),  38. 


;j.  ZINC. 

Zn  65.4. 

The  natural  compounds  of  the  heavy  metals  liave  generally  a high  spe- 
cific gravity,  frequently  jiossess  metallic  luster,  usually  occur  in  the  older 
crystalline  rocks  in  veins,  and  are  termed  ores.  The  most  important  zinc 
ores  are  the  carbonate,  ZnCO,,  the  silicate,  and  sphalerite  or  blende,  ZnS. 
The  iirincipal  sources  of  these  ores  are  in  vSilesia,  hingland,  Belgium, 
Poland,  the  United  ^States,  Prance,  and  Spain.  To  get  the  metal  the 
carbonate  or  sulphide  is  converted  into  oxide  by  roasting  it  in  the  air;  the 
product  is  then  mixed  with  carbon  and  ignited  in  cylindrical  clay  tubes. 
In  this  manner  the  oxide  is  reduced  : 

ZnO  -L  C — Zn  CO, 

and  the  liberated  zinc  distilled  off.  The  receivers  contain  the  fused, 
compact  zinc  and  a gray,  pulverulent  mass,  called  zinc-dust,  which  con- 
sists of  a mixture  of  zinc  oxide  with  finely  divided  metal.  This  material 
is  used  in  laboratories  as  a strong  reducing  agent.  See  Mylius  and  Fromm 
on  the  purification  of  zinc,  Z.  f.  anorg.  Ch.  g (1895),  144. 

Metallic  zinc  has  a bluish-white  color,  and  exhibits  rough,  crystalline 
fracture;  its  specific  gravity  equals  7-7.  i . At  ordinary  temperatures  it 
is  brittle  and  can  be  pulverized;  at  100-150°  it  is  malleable  and  can  be 
rolled  into  thin  leaves  and  drawn  out  into  wire.  At  200°  it  becomes 
brittle  again  and  may  be  easily  broken.  It  melts  at  420°  and  distils 
at  about  950°. 

It  becomes  coated  with  a thin  layer  of  basic  carbonate  in  moist  air. 
Heated  in  the  air  it  burns  to  zinc  oxide  with  a very  intense,  bluish-white 
light.  Compact  zinc  will  only  decompose  water  at  a red  heat ; zinc-dust, 
however,  acts  at  ordinary  temperatures.  Perfectly  pure  zinc  is  only 
slowly  attacked  by  dilute  acids  at  the  ordinary  temperature;  it  dissolves 
by  boiling  with  potassium  or  sodium  hydroxide,  as  well  as  ammonia, 
with  liberation  of  hydrogen  : 

Zn  4-  2NaOII  = Zn(ONa)2  + Hj. 

Owing  to  its  slight  alteration  in  the  air  zinc  meets  with  extensive  appli- 
cation as  sheet-zinc  for  coating  statues  and  in  architectural  adornment, 
and  in  galvanizing  sheet-iron.  It  also  forms  an  important  constituent 
of  many  valuable  alloys,  such  as  brass  and  argentan  (see  these). 

Zinc  Hydroxide,  ZnfOH)^,  is  preci])itated  as  a white,  amorphous 
powder,  from  acpieous  solutions  of  zinc  salts,  by  alkalies,  and  is  soluble  in 
excess  of  the  reagent.  When  heated  it  deconq)oses  into  water  and  zinc 
oxide. 

Zinc  Oxide,  ZnO,  is  usuallv  jirepared  by  igniting  the  precipitated 
basic  carbonate,  and,  as  zinc  iviiile,  is  employed  as  a stable  white  ])aint. 
'I'he  oxide  obtained  by  burning  the  metal  is  a white,  voluminous,  floccu- 


2INC.  3^9 

lent  mass,  called  Flores  zi/ici,  or  Lana philosopJiica.  When  zinc  oxide  is 
heated  it  acquires  a yellow  color,  which  disappears  on  cooling. 

Zinc  oxide  occurs  in  nature  as  zincite^  colored  by  impurities. 

Zinc  Chloride,  ZnCl2  {Zincutn  chloratiini),  anhydrous,  is  obtained  by 
heating  zinc  in  a stream  of  chlorine,  by  the  evaporation  of  the  solution 
of  zinc  in  hydrochloric  acid,  and  by  the  distillation  of  zinc  sulphate  with 
calcium  chloride.  It  forms  a white,  deliquescent  mass,  which  fuses  when 
heated  and  distils  at  about  730°.  Molten  zinc  chloride  is  a transparent, 
very  mobile,  strongly  refracting  liquid,  attracting  moisture  more  readily 
than  phosphoric  anhydride  [Lorenz,  Z.  f.  anorg.  Ch.  10  (1895),  82].  It 
crystallizes  from  concentrated  hydrochloric  acid  solution  with  one  mole- 
cule of  water  in  deliquescent  octahedra.  When  the  aqueous  solution  of 
zinc  chloride  is  evaporated  it  partly  decomposes  (like  magnesium  chlo- 
ride) into  zinc  oxide  and  hydrochloric  acid.  When  the  concentrated 
zinc  chloride  is  mixed  with  zinc  oxide,  a plastic  mass  is  obtained,  which 
hardens  rapidly;  a mixture  of  magnesium  chloride  and  oxide  does  the 
same.  In  both  instances  the  hardening  depends  upon  the  formation  of 

Cl  . . 

basic  oxychlorides,  e.  g.,  Zn<Qp^  Zinc  chloride  forms  deliquescent 

double  salts  with  the  alkaline  chlorides,  e.  g.,  ZnCl2.2KCl.  With 
ammonia  it  yields  various  compounds,  of  which  ZnCl2.NH3  is  character- 
ized by  great  stability. 

Zinc  Sulphate,  ZnSO^,  is  obtained  by  dissolving  zinc  in  sulphuric 
acid.  It  is  prepared  upon  a large  scale  by  carefully  roasting  zinc- 
blende  (ZnS) ; the  zinc  sulphate  is  extracted  by  water.  It  crystallizes 
at  ordinary  temperatures  from  aqueous  solutions  with  seven  molecules  of 
water-(zinc  or  white  vitriol,  Zincum  sulphuricum')  in  rhombic  crystals,  re- 
sembling those  of  magnesium  sulphate.  It  affords  double  salts  with  the 
alkaline  sulphates;  these  contain  six  molecules  of  water  (p.  315). 

Zinc  Carbonate,  ZnCOg,  occurs  native  as  smithsonite  in  hexagonal 
crystals,  isomorphous  with  those  of  calcite.  Sodium  carbonate  precipi- 
tates a mixture  of  ZnCOg.  H2O,  which  rapidly  becomes  crystalline,  and 
the  basic  carbonate  2ZnC03. 3Zn(OH)2. H2O. 

Zinc  Sulphide,  ZnS,  is  zinc-blende  or  sphalerite,  usually  colored 
brown  by  ferric  oxide  or  other  admixtures.  Ammonium  sulphide  precipi- 
tates it  as  a white  compound  from  zinc  solutions.  Although  fused  zinc 
reacts  with  difficulty  with  sulphur,  zinc-dust  combines  with  the  latter  in 
powdered  form  quite  readily,  and  if  the  mixture  be  heated  or  struck  with 
a hammer  the  union  is  accompanied  by  an  explosion.  Zinc  sulphide  is 
insoluble  in  water,  but  is  readily  dissolved  by  dilute  acids,  excepting  acetic 
acid;  therefore  it  may  be  precipitated  by  hydrogen  sulphide  from  zinc 
acetate  solutions.  This  reaction  serves  to  separate  zinc  from  other  metals. 
The  sulphide  is  also  applied  as  a white  pigment ; especially  a mixture  of 
it  with  barium  sulphate  (lithopone — made  from  zinc  sulphate  and  barium 
sulphide). 

Zinc  Silicate,  Zn2SiO^-(-  H2O,  occurs  in  rhombic  crystals  as  calamine. 
In  the  analytical  way  zinc  is  characterized  by  the  fact  that  its  hydroxide 
dissolves  both  in  caustic  soda  and  potash  as  well  as  in  ammonia;  hydro- 
gen sulphide  precipitates  white  zinc  sulphide  from  these  solutions. 


320 


INORGANIC  CHEMISTRY. 


4.  CADMIUM. 

Ccl  = 1 12. 

Cadmium  very  often  accompanies  zinc  in  its  ores.  As  much  as  5 per 
cent,  of  this  metal  is  present  in  the  .Silesian  zinc  ores;  it  was  first  dis- 
covered in  them  in  181  7 by  Hermann  ; shortly  after  Stromeyer  recognized 
it  as  a new  element.  Jleing  more  volatile  than  zinc,  in  olitaining  the 
latter  it  distils  off  first,  and  may  be  easily  separated  from  the  first  por- 
tions of  the  distillate.  It  is  a white,  tenacious,  and  rather  soft  metal,  of 
specific  gravity  8.6.  It  fuses  at  320°  and  l)oils  at  770°.  It  does  not  alter 
much  in  the  air.  Heated,  it  burns  with  the  separation  of  a brown  smoke 
of  cadmium  oxide.  It  dissolves  with  difficulty  in  dilute  hydrochloric  and 
sulphuric  acids,  but  readily  in  nitric  acid.  Zinc  throws  out  the  metal 
from  solutions  of  the  soluble  cadmium  salts. 

St.  Claire  Deville  found  the  specific  gravity  of  cadmium  vapors  at  1040°  and  II. 
Biltz  at  about  1700°  to  approximate  closely  112  (0.^;=r32).  'i’herefore,  the  molecular 
weight  of  cadmium  is  112.  Since  the  atomic  weight  of  cadmium  (determined  from  its 
specific  heat)  is  also  1 12,  it  follows  that  the  gas  molecule  of  cadmium  consists  of  but 
atom.  We  know  that  the  molecules  of  other  elements  in  the  gaseous  state  are  composed 
of  two  or  more  atoms  (O.^,  N2,  P4).  Cadmium,  therefore,  forms  an  exception  to  this  rule. 
This  is  also  true  of  mercury,  zinc  and  perhaps,  too,  of  other  bivalent  metals.  The.se 
relations  remind  us  of  the  behavior  of  the  hydrocarbon  residues  (radicals)  ; while  the 
bivalent  or  quadrivalent  groups,  <?.  , ethylene.  C.^H^,  and  acetylene,  exist  in  free 

condition,  the  univalent  groups  (as  CH,  and  CNj  cannot  appear  free,  but  double  them- 
selves when  separated  from  their  compounds. 

Cadmium  Hydroxide,  Cd(0H)2,  is  precipitated  as  a white  powder, 
from  the  soluble  cadmium  salts,  by  the  alkalies;  it  is  insoluble  in  sodium 
and  potassium  hydroxides,  but  dissolves  readily  in  ammonium  hydroxide. 

Cadmium  Oxide,  CdO,  is  prepared  by  igniting  the  nitrate.  It  is  a 
brownish-black  ])owder,  consisting  of  microscopic  octahedra.  When  ob- 
tained by  heating  the  carbonate  or  hydroxide,  it  is  a brown,  amorphous 
powder. 

Cadmous  Hydroxide  CdOH,  and  its  oxide,  Cd^O,  have  been  prepared.  The  first 
is  a grayish- white  compound,  while  the  second  is  composed  of  yellow,  translucent  crystals. 
It  is  obtained  on  heating  the  hydroxide  to  a temperature  at  which  sulphuric  acid  gives 
off  dense  white  fumes.  The  hydroxide  is  a reducing  agent,  yielding  hydrogen  with 
hydrochloric  acid  and  oxides  of  nitrogen  with  nitric  acid.  The  oxide  conducts  itself 
similarly.  These  compounds  are  of  interest,  as  they  foreshadow  the  tendency  toward 
the  formation  of  lower  oxides,  so  strongly  shown  by  mercury  {^Am.  Cheni.  Jr.,  12,  493  . 

Cadmium  Chloride,  CdCl2,  crystallizes  from  aqueous  solution,  with 
two  molecules  of  water,  and  may  be  dried  without  decomposition.  The 
anhydrous  salt  melts  at  540°  and  sublimes  in  scales. 

Cadmium  Iodide,  Cdl2,  is  obtained  by  the  direct  action  of  iodine 
upon  metallic  cadmium  in  the  j)resence  of  water.  It  crystallizes  from 
the  latter  in  hexagonal  tal)les.  It  is  used  in  photography. 

Cadmium  Sulphate,  CdSOj,  crystallizes  at  — 20°,  like  the  sulphates 
of  zinc  and  magnesium  at  the  ordinary  temperature,  with  seven  mole- 
cules of  water,  but  at  the  ordinary  temj)erature  with  -I  molecules  of 
water  ; the  crystals  effloresce  in  the  air.  It,  however,  forms  double  salts 


CADMIUM. 


321 


with  the  sulphates  of  the  alkali  metals,  e.  g.,  CdSO^  . K.^SO^  -f  dH/) ; 
these  are  perfectly  analogous  to  thv/Se  of  zinc  and  magnesium,  and  iso- 
morphous  with  them  (p.  319). 

Cadmium  Sulphide,  CdS,  occurs  native  as  greenockite,  in  yellow 
hexagonal  prisms.  Hydrogen  sulphide  precijutates  it  from  cadmium  salt 
solutions  as  a yedow  powder,  insoluble  in  dilute  acids.  It  is  employed 
as  a pigment. 

Almost  all  the  alloys  of  cadmium  have  a low  melting  point.  Freshly 
prepared  cadmium  amalgam  is  a white  ])lastic  mass,  which  soon  becomes 
hard ; it  was  formerly  used  for  filling  teeth. 


The  magnitude  of  the  atomic  weight  of  Mercury  would  place  the  lat- 
ter in  the  group  of  zinc  and  cadmium.  The  relationship  of  these  three 
heavy  metals  occupying  a similar  position  in  the  three  great  periods  are, 
from  a physical  point  of  view,  distinguished  among  the  heterologous  mem- 
bers by  their  ready  fusibility  and  volatility,  which  nearly  reach  a maxi- 
mum in  them.  In  the  homologous  series,  zinc,  cadmium,  mercury,  these 
properties  increase  with  rising  atomic  weight  (just  as  with  the  metals  of 
the  potassium  group,  p.  272)  : 


Zn 

Cd 

Hff 

Atomic  weight,  

65-4  ' 

I12 

200.3 

Melting  point, 

420° 

320° 

-39-5° 

Boiling  point, 

950° 

770° 

357° 

Specific  gravity, 

7.1 

8.6 

13-5 

Atomic  volume, 

9.1 

13 

14.7 

The  gradation  in  the  heat  of  formation  of  their  compounds  (p.  327) 
clearly  indicates  that  mercury  must  be  arranged  in  a group  with  cad- 
mium and  zinc. 

Like  zinc  and  cadmium,  it  yields  compounds  of  the  form  HgX^,  in 
which  it  a],)pears  bivalent.  These  derivatives  are,  in  many  respects,  simi- 
lar to  the  corresponding  compounds  of  zinc  and  cadmium.  Thus,  mer- 
curic sulphate  forms  double  salts  with  the  alkaline  sulphates,  which  crys- 
tallize with  six  molecules  of  water  (HgS04.K2S0^  -f  6H2O),  and  are  iso- 
morphous  with  the  double  sulphates  of  the  magnesium  group  (p.  315). 
The  similarity,  however,  limits  itself  to  few  compounds.  Since  the 
properties  of  each  group  sustain  a slight  change  by  virtue  of  the  increas- 
ing atomic  weight,  we  are  not  surprised  to  observe  this  to  be  very  evident 
in  the  case  of  mercury  (with  the  high  atomic  weight  200.3),  especially  as 
the  middle  ftransition)  member  of  the  third  great  period  is  not  known 
(p.  245).  Mercury  differs  essentially  from  zinc  and  cadmium  in  that,  in 

II 

addition  to  the  compounds  of  the  form  HgX2  (mercuric  compounds),  it 

' I 

is  also  capable  of  yielding  those  of  the  form  HgX  (mercurous  compounds^ 
in  which  it  seems  to  be  univalent.  Here  we  meet  an  instance,  frequently 


322 


in()R(;anic  chemis'I'rn  . 


observed,  in  wliich  one  and  the  same  metal  (as  witli  tlic  most  metalloids) 
is  capable  of  forming  compounds  of  two  or  more  forms,  which  are  to  be 
referred  to  a different  valence  of  the  metal  ; and  it  often  haj)i)ens  that 
the  derivatives  of  a metal,  ap[)earing  in  different  forms  or  types,  are  fre- 
quently more  essentially  distinguished  from  one  another  than  the  com- 
pounds of  different  elements  having  the  same  tvpe.  Thus,  the  mercuric 

II 

compounds  (IlgX.J  are  similar  to  those  of  zinc  and  cadmium,  after  the 

I 

same  form,  while  the  mercurous  com})ounds  (HgX)  exhibit  great  resem- 
I I 

blance  to  the  cuprous  (CuX)  and  silver  (AgX)  compounds,  constituted  • 
according  to  a similar  tyj)e. 

It  shows  that  the  similarity  of  the  compounds  is  influenced  not  only  by 
the  nature  of  the  metals,  but  frequently,  to  a marked  degree,  by  the 
forms  or  types  according  to  which  they  are  constituted  (p,  330). 

Heretofore  it  has  been  generally  customary  to  assume  the  presence  of  two  atoms  of 
mercury,  united  to  a bivalent  group,  in  the  molecule  of  the  mercurous  derivatives  ; hence 
the  appearance  of  a univalent  character  : 

llg-Cl  Hg-NOs 

I'lg-Cl  Hg-NO.,. 

Chloride.  Nitrate. 

The  most  recent  researches,  however,  show  that  the.se  compounds  can  be  represented 
by  the  simple  formulas  HgCl,  HgN03  ; see  pp.  325,  332. 


5.  MERCURY. 

Hg  = 200.3. 

Mercury  {Hydrargyrum,  i.  e.,  water-silver')  occurs  in  nature  principally 
as  the  sulphide,  cinnabar ; more  rarely  native  in  the  form  of  little  drops 
scattered  through  rocks.  Its  most  important  localities  are  Almaden,  in 
Spain;  New  Almaden,  in  California;  Idria,  in  Austria;  Tuscany,  in 
Italy;  Mexico  ; Peru;  China;  Japan.  Considerable  quantities  have  been 
found  also  in  Russia. 

The  metallurgical  separation  of  mercury  is  very  simple.  Cinnabar  is 
roasted  in  reverberatory  furnaces,  whereby  the  sulphur  burns  to  dioxide, 
and  the  mercury  vapors  are  condensed  in  large  chambers  with  the  aid  of 
water.  There  is  a regular  and  gradual  return  to  the  old  method  of  dis- 
tilling cinnabar  with  lime  or  iron  from  iron  retorts.  Commercial  mer- 
cury usually  contains  a slight  quantity  of  other  metals  dissolved  in  it. 
h'or  its  purification,  it  is  poured  in  a thin  stream  into  a deep  layer  of  dilute 
nitric  acid  or  of  ferric  chloride  solution,  or  it  is  shaken  with  a solution 
of  i)otassium  dichromate  and  dilute  sulifliuric  acid,  by  which  the  admixed 
metals  are  more  easily  di.ssolved  than  the  mercury.  Finally,  the  metal  is 
passed  through  chamois  skin  and  distilled  (best  in  a vacuum).  Ber.  12 
(1879),  204,  576. 

Mercury  is  the  only  metal  which  is  liquid  at  ordinary  temperatures. 
At  0°  its  s|)ecific  gravity  equals  13.596  and  at  15°,  13.559  ; it  solidi- 
fies at  — 39-5°,  and  crystallizes  in  regular  octahedra  ; it  evaporates  some- 


MERCUROUS  COMPOUNDS. 


323 


what  at  medium  temperatures,  and  boils  at  357°.  Its  vapors  are  very 
poisonous.  The  specific  gravity  of  the  vai)or  of  mercury  is  200.3, 
(O2  = 32).  Therefore,  the  molecular  weight  of  the  metal  is  200.3, 
its  atomic  weight  is  also  200.3,  molecule,  like  that  of  cadmium  and 
zinc,  is  composed  of  only  one  atom.  At  ordinary  temperatures,  mercury 
is  not  altered  by  exposure  to  the  air  ; near  the  boiling  point,  however,  it 
gradually  oxidizes  to  red  mercuric  oxide.  Hydrochloric  and  cold  sulphuric 
acids  do  not  act  upon  mercury  ; hot  sulphuric  acid  converts  it  into  mer- 
curic sulphate,  with  evolution  of  sulphur  dioxide.  Even  dilute  nitric  acid 
will  readily  dissolve  it.  It  combines  with  the  halogens  and  sulphur  at 
ordinary  temperatures. 

Mercury  dissolves  almost  all  metals  (not  iron)  forming  amalgams.  It 
unites  with  potassium  and  sodium  upon  gentle  warming,  with  production 
of  heat  and  light.  When  the  quantity  of  potassium  and  sodium  exceeds 
2 per  cent.,  the  alloy  is  solid  and  crystalline;  by  less  amount  it  remains 
liquid.  [Compare  Z.  f.  phys.  Ch.  29  (1899),  119.]  Tin  amalgam  is 
employed  for  coating  mirrors. 

Sodium  amalgam  may  be  made  by  placing  rather  large  pieces  of  sodium  upon  the 
bottom  of  an  iron  crucible,  covering  them  with  wire  gauze,  so  that  they  cannot  rise  to  the 
surface,  and  pouring  mercury  upon  them.  With  large  quantities  the  reaction  begins  cf 
itself.  Finally,  heat  must  be  applied  while  stirring. 


Mercury  forms  two  series  of  compounds,  mercurous  and  mercuric.  The 
first  are  analogous  to  the  cuprous,  and  have  the  form  HgX.  In  them 
mercury  appears  to  be  univalent.  In  many  respects  the  mercurous  com- 
pounds are  similar  to  the  cuprous  and  silver  derivatives.  The  halogen 
compounds  are  insoluble,  and  darken  on  exposure  to  light. 

In  the  mercuric  derivatives,  HgX2,  mercury  is  bivalent,  and  is  very 
much  like  zinc  and  cadmium.  Mercuric  compounds  almost  always  form, 
if  the  substance  reacting  with  the  mercury  is  in  excess;  when  the  oppo- 
site is  the  case,  mercurous  salts  result.  Mercuric  derivatives,  by  the 
addition  of  mercury,  pass  into  the  mercurous,  e.  g., 

Hg(NO,y+  Hg  = Hg,(NO,),. 

This  change  may  also  be  effected  by  reducing  agents  (sulphurous  acid, 
phosphorous  acid  and  stannous  chloride).  Oxidizing  agents  convert  the 
mercurous  into  mercuric  compounds. 


MERCUROUS  COMPOUNDS. 

Mercurous  Chloride,  HgCl  or  Hg2Cl2,  calomel  {Hydrargyrum  chlor- 
atum),  is  an  amorphous  white  precipitate,  produced  by  the  addition  of 
hydrochloric  acid  or  soluble  chlorides  to  the  solution  of  mercurous  salts. 
It  is  generally  formed  by  the  sublimation  of  mercuric  chloride  with  mer- 
cury ; or  a mixture  of  mercuric  sulphate,  mercury  and  sodium  chloride 
is  sublimed  : 

IlgSO,  -f  2NaCl  -f  Hg  = Na2SO,  + 2HgCl. 


324 


INORGANIC  CHEMISTRY. 


It  then  forms  a radiating,  crystalline  mass  (cjuadratic  prisms)  of  sjiecific 
gravity  7.2.  Calomel  is  insoluble  in  water,  in  alcohol,  and  in  dilute  acids  ; 
it  gradually  decomi)oses  when  exi)osed  to  the  light,  with  separation  of 
mercury.  When  heated,  it  sublimes  without  fusing,  foiling  hydro- 
chloric acid  and  hot  concentrated  solutions  of  calcium  chloride  and 
alkaline  chlorides  decompose  it  into  mercuric  chloride  and  free  mercury: 

2lIgC]  = HgCl,  + Ilg. 

When  ammonium  hydroxide  is  j^oured  over  calomel,  it  blackens  (hence 
the  name  calomel,  from  xak(>/xeX(x',  beautiful  black);  it  is  not  surely 
known  what  chemical  change  occurs  heie. 

The  vapor  density  of  calomel  vapors  at  400°  (first  determined  by  Mitscherlich,  and  con- 
firmed by  Deville  and  Troost,  Rieth  and  (Idling)  is  235  (<  >2  = 32)  ; which  corresponds 
to  the  formula  HgCl.  As  formerly  supposed,  and  as  recently  demonstrated  by  V.  Meyer 
and  Harris  in  opposition  to  Fileti,  calomel  breaks  down  completely  into  mercury  and 
mercuric  chloride  [Her.  28  (1895),  364].  This  mixture  has  the  same  vapor  density  that 
undecomposed  calomel  would  show  : 

iigci  + Hgci  = ng+  iigcb- 

I vol.  I vol.  I vol.  I vol. 

The  question,  whether  the  mercurous  compounds  contain  one  or  two  atoms  of  mer- 
cury, whether,  for  exam[)le,  the  formula  Hg.^CI.^  or  llgCl  properly  belongs  to  calomel, 
can,  therefore,  not  be  decided  by  the  determination  of  its  vapor  density  alone.  It  must 
also  be  ascertained  of  what  the  vapor  consists — and  this  proof  free  from  objection  has  not 
yet  been  given.  From  electro-chemical  experiments  and  from  the  analogy  to  cuprous 
and  silver  chlorides  it  would  seem  very  probable  that  mercurous  chloride  has  the  simple 
formula  HgCl  [Z.  f.  anorg.  C h.  9 (18951,  442]. 

Mercurous  Nitride,  Ngllg  or  Hg.2(N3l2,  is  formed  by  adding  mercurous  nitrate  to 
solutions  of  hydrazoic  acid  or  its  alkali  salts.  It  is  insoluble  in  water.  It  consists  of 
microcrystalline  needles,  becoming  yellow  in  the  light  and  blackened  by  ammonia,  just 
like  calomel.  When  heated  or  struck  it  explodes  with  great  violence.  Its  separation 
into  its  elements  is  accompanied  with  a brilliant  blue  light  (p.  312). 


Mercurous  Iodide,  Hgl  or  Hg2l2»  is  prepared  by  rubbing  together 
8 parts  of  mercury  with  5 parts  of  iodine,  or  by  precipitating  mercurous 
nitrate  with  potassium  iodide.  It  is  a yellowish-green  powder,  insoluble 
in  water  and  in  alcohol.  Light  changes  it  to  mercuric  iotiide  and  mercury. 
Aqueous  solutions  of  potassium  iodide  have  the  same  effect. 

Mercurous  Oxide,  Hg^O,  is  a black  mass,  and  is  formed  by  the 
action  of  potassium  or  sodium  hydroxide  upon  mercurous  salts.  In  the 
light  or  at  100°  it  decomposes  into  mercuric  oxide  and  mercury. 

Mercurous  Nitrate,  HgNOg  or  Hg2(N03\,  is  produced  by  allowing 
dilute  nitric  acid  to  act  upon  excess  of  mercury  in  the  cold.  It  crystal- 
lizes with  one  molecule  of  water  in  large  monoclinic  tables.  It  dissolves 
readily  in  water  acidulated  with  nitric  acid  ; pure  water  partly  decom- 
poses it  with  the  separation  of  a yellow  basic  salt  of  the  composition 


„ OH 

The  nitric  acid  solution  of  mercurous  nitrate  oxidizes  when  exposed  to 
the  air,  and  gradually  becomes  mercuric  nitrate;  this  may  be  jirevented 
by  adding  metallic  mercury  to  the  solution,  whereby  the  resultant  mer- 
curic salt  is  again  changed  to  the  mercurous  state: 


Hg.N03)2  f Ilg  = Hg,(N03). 


MERCURIC  COMPOUNDS. 


325 


Mercurous  Sulphate,  Hg2(SOJ,  results  when  an  excess  of  mercury 
is  heated  gently  with  concentrated  sulphuric  acid  ; it  separates  as  a crystal- 
line precipitate,  difficultly  soluble  in  water,  if  sulphuric  acid  be  added  to 
a mercurous  nitrate  solution.  It  fuses  upon  ap})lication  of  heat,  and  de- 
composes into  sulphur  dioxide,  oxygen  and  mercury. 

Mercurous  Sulphide,  HggS,  is  not  definitely  known.  A black  com- 
pound is  produced  in  dilute  solutions  of  mercurous  nitrate  by  hydrogen 
sulphide  or  alkaline  sulphides.  It  contains  mercury  and  mercuric  sul- 
phide : 

Hg^S  =:  Hg  + HgS. 


MERCURIC  COMPOUNDS. 

Mercuric  Chloride,  HgClj,  corrosive  sublimate  (^Hydrargyrum  bichlo- 
ratum),  is  produced  when  mercuric  oxide  is  dissolved  in  hydrochloricacid, 
or  metallic  mercury  in  aqua  regia.  It  is  obtained  on  a large  scale  by  the 
sublimation  of  a mixture  of  mercuric  sulphate  with  sodium  chloride: 

HgSO,  + 2NaCl  = HgCb  + Na^SO,. 

It  crystallizes  from  water  in  rhombic  prisms,  and  dissolves  at  medium 
temperatures  in  15  parts  of  water,  at  100°  in  2 parts;  it  is  still  more 
soluble  in  alcohol.  Its  specific  gravity  is  5.4.  It  fuses  at  265°  and 
boils  at  307°,  Its  critical  pressure  is  about  420  mm.  (p.  229).  The 
vai)or  density  corresponds  to  the  molecular  formula  HgClj  (=  271.2). 

Reducing  substances,  like  sulphur  di'<xide  and  stannous  chloride, 
change  it  to  insoluble  mercurous  chloride: 

2HgCl2  -f  SO2  + 2H2O  = Hg2Cl2  + H2S0^  + 2HCI. 

Stannous  chloride  first  precipitates  mercurous  chloride: 

2HgC]2  + SnClj  = Hg2Cl2  + SnCh, 

which  is  afterward  reduced,  by  the  excess  of  stannous  chloride,  to  me- 
tallic mercury  : 

iig2a2  + Sncq  = 2Hg  -f  Sncq. 

Mercuric  chloride  is  greatly  inclined  to  form  double  salts  with  metallic 
chlorides,  e.  g.,  HgCl2.KCl  -j-  H2O.  These  may  be  regarded  as  the 
salts  of  a hydrochlormercuric  acid,  H.HgClgfp.  271).  The  acidity  of 
an  aqueous  solution  of  mercuric  chloride  is  neutralized  by  the  addition 
of  sodium  chloride,  which  would  indicate  the  formation  of  such  salts. 
When  ammonium  hydroxide  is  added  to  its  solution,  a heavy  white  pre- 
cipitate is  produced.  Its  composition  varies  with  the  concentration,  the 
temperature,  and  the  quantities  of  mercuric  chloride  and  ammonia. 
White  precipitate  (^Hydrargyrum  prcecipitatum  alinin')  has  the  formula 
HgClNHj.  This  compound  may  be  regarded  as  mercuric  chloride,  in 


326 


INORGANIC  CHKMISTRY. 


which  one  atom  of  chlorine  is  replaced  Ijy  the  amido-gronp,  NH,  : 

, and  it  has  been  called  amido-mcrcury  chloride.  Or  it  may 

be  derived  from  ammonium  chloride  by  the  substitution  of  a mercury 
atom  for  two  hydrogen  atoms,  NHglljCl,  mercuramnumium  chloride. 
Similar  mercuri-  and  mercuro-amimmium  derivatives  are  numerous. 

Mercuric  Iodide,  Hgl2  (^Hydrargyrum  hiiodafum),  is  formed  by  the 
direct  union  of  mercury  with  iodine.  When  i)otassium  iodide  is  added 
to  a solution  of  mercuric  chloride,  mercuric  iodide  separates  as  a yellow 
precipitate,  which  immediately  becomes  red.  It  is  readily  soluble  in 
mercuric  chloride  and  ])otassium  iodide  solutions;  it  crystallizes  from 
alcohol  in  bright  red  cpiadratic  pyramids.  Upon  heating  dry  mercuric 
iodide  to  150°,  it  suddenly  becomes  yellow,  and  fuses  at  223°  to  a red 
liquid;  a i)ortion  subliming  before  this  in  yellow,  shining,  rhombic 
needles.  On  touching  these  with  some  solid,  they  become  red,  with 
sei)aration  of  heat,  and  are  changed  into  an  aggregate  of  quadratic 
pyramids.  Mercuric  iodide  is  therefore  dimorphous.  It  crystallizes 
(depending  on  the  temperature),  in  either  yellow  or  red  forms,  from 
methylene  iodide,  01^^12,  of  which  100  jjarts  dissolve  16.6  parts  of  mer- 
curic iodide  at  100°. 

The  great  resistance  shown  by  all  the  mercury  halides  toward  concen- 
trated sulphuric  acid  is  noteworthy.  Even  in  the  heat  the  halogen  is  very 
slowly  ex])elled.  They  are  also  very  stable  toward  caustic  soda  and  pot- 
ash. This  is  in  harmony  with  the  tendency  of  mercury  to  yield  deriva- 
tives which  are  only  slightly  dissociated. 

Mercuric  Cyanide,  Hg(CN)2,  is  obtained  by  dissolving  yellow  mercuric  oxide  in 
aqueous  prussic  acid  or  by  heating  Prussian  blue  with  mercuric  oxide  and  water.  It  crys- 
tallizes from  water  in  white  quadratic  prisms.  It  is  also  soluble  in  alcohol  and  can  be 
extracted  from  its  aqueous  solution  by  shaking  with  ether.  It  breaks  down  on  heating 
into  mercury  and  cyanogen  gas  (p.  236).  The  mercury  and  cyanogen  group  seem  to  be 
differently  combined  from  other  metals  and  cyanogen,  e.  g.,  potassium  cyanide.  Mercuric 
cyanide  has  a very  low  degree  of  dissociation.  Silver  nitrate  does  not  precipitate  silver 
cyanide  from  its  aqueous  solution,  nor  does  an  alkaline  hydroxide  throw  out  mercuric 
oxide.  Hydrogen  sulphide  and  alkaline  sulphides  decompose  it  with  formation  of  mer- 
cury sulphide  and  hydrocyanic  acid  or  alkaline  cyanide.  Unlike  potassium  cyanide  it  is 
very  stable  toward  acids. 

Mercuric  Oxide,  HgO,  is  obtained  by  the  prolonged  heating  of 
metallic  mercury  near  the  boiling  point  in  the  air,  orUy  the  ignition  of 
mercurous  or  mercuric  nitrate  and  mercury.  It  forms  a red,  crystalline 
powder  (^Hydrargyrum  oxydatum  rubruui),  of  specific  gravity  1 1.2.  When 
sodium  hydroxide  is  added  to  a solution  of  mercuric  chloride,  mercuric 
oxide  separates  as  a yellow,  amorjihous  itrecijn'tate  {Hydr.  oxyd.  flavum  via 
hu?nida  paratuni).  Both  modifications  become  black  when  heated,  but 
change  to  a yellowish-red  on  cooling.  Mercuric  oxide  decomposes  into 
mercury  and  oxygen  at  about  400°. 

Mercuric  oxide  combines  directly  with  ammonia,  to  form  the  compound 
2HgU.  NII3,  which  explodes  with  violence  when  heated. 

Mercuric  Nitrate,  Ilg^NO^b^-  It  is  difficult  to  obtain  this  salt  pure, 
because  it  is  inclined  to  form  basic  com})ounds.  A solution  of  it  may  be 


MERCURIC  COMPOUNDS. 


327 


made  by  dissolving  mercury  or  mercuric  oxide  in  an  excess  of  hot  nitric 
acid.  On  diluting  the  solution  with  water  the  basic  salt,  Hg(N03)2.  2HgO 
-j-  HjO,  separates,  and  this  may  be  converted  into  pure  mercuric  oxide 
by  boiling  with  water. 

Mercuric  Sulphate,  HgSO^,  is  produced  by  digesting  mercury  or 
its  oxide  with  an  excess  of  concentrated  sulphuric  acid.  It  forms  a white, 
crystalline,  insoluble  mass,  which  becomes  yellow  on  heating.  It  yields 
the  hydrate  HgSO^  -f-  H^O  with  a little  water,  but  much  of  the  latter 
decomposes  it  into  sulphuric  acid  and  the  yellow  insoluble  basic  salt, 
HgSO^.  2HgO  ( Turpcthum  mme7'ale,  Turpeth  mineral). 

Mercuric  sulphate  forms  double  salts  wjth  the  alkaline  sulphates,  e.  g., 
HgSO^.K2SO^  6H2O ; these  are  isomorphous  with  the  corresponding 
double  salts  of  the  magnesium  grouj)  (p.  315). 

Mercuric  Sulphide,  HgS,  0(  curs  in  nature  as  cinnabar,  in  radiating 
crystalline  masses,  or  in  hexagonal  prisms  of  dark-red  color.  It  is  obtained 
artificially  by  rubbing  together  mercury  and  flowers  of  sulphur  with  water, 
or  it  is  produced  as  a black  microcrystalline  precipitate  by  the  precipita- 
tion of  a solution  of  a mercuric  salt  with  hydrogen  sulphide.  If  the  black 
sul])hide  be  heated  with  exclusion  of  air  it  sublimes  as  a dark-red  mass  of 
radiating  crystalline  structure,  and  is  perfectly  similar  to  natural  cinnabar. 
A similar  conversion  of  the  black  modification  into  the  red  is  effected  by 
continued  heating  of  the  same  to  50°  with  a solution  of  potassium  or 
ammoniumsulphide.  The  red  mercury  sulphide  thus  obtained  is  employed 
as  artificial  cinnabar  in  painting. 


The  mercury  compounds  can  be  readily  recognized  by  the  following 
reactions.  When  heated  with  dry  sodium  carbonate  in  a small  test- 
tube,  mercury  escapes,  and  condenses  u])on  the  side  in  metallic  drops. 
Tin,  copper  and  zinc  throw  out  metallic  mercury  from  its  solutions. 
If  a piece  of  sheet  copper  be  dipped  into  the  same,  mercury  is  de- 
posited as  a gray  coating,  which  on  being  rubbed  acquires  a metal- 
lic luster.  When  a dry  mixture  of  a mercuric  salt,  lime  and  potassium 
iodide  is  heated  yellow  (or  red)  mercuric  iodide  sublimes.  This  is 
the  only  iodide  not  decomposed  by  caustic  lime.  The  mercurous  com- 
pounds are  distinguished  from  the  mercuric  by  their  precipitation  by 
hvdrochloric  acid. 


The  heats  of  formation  of  the  chlorides  of  the  seven  metals  of  the  second  group  are 
as  follows : 


(Mg,Cl2)  = 15 1 

(Ca,Cb)  170  (Zn,Cl2)  =--  97 

(Sr,Cl2)  = 185  (Cd,Ch)  = 93 

(Ba,Cl2)  = 195  (llg,('l,)  -54. 


I'he  heats  of  formation  in  the  first  sub-class  increase  with  the  rise  in  the  atomic  weights, 
while  the  opposite  is  observed  in  the  second  sub-class.  The  heats  of  formation  of  other 
derivatives  of  the  two  groups  manifest  similar  relations. 


328 


INuKGANIC  CHEMISTRY. 


Comparing  these  numbers  with  the  quantity  of  heat  which  is  disengaged  in  the  forma- 
tion of  aqueous  hydrochloric  acid  : 

II,C1  = 22;  ir,Cl,Aq  39.3, 

we  find  explained  the  behavior  of  the  metals  toward  this  acid.  All  metals  liberating  a 
greater  (juantity  of  heat  than  39.3  Cal.  in  the  formation  of  their  chlorides  (calculated  for 
I equivalent  of  metal)  are  in  condition  to  decompose  the  dilute  acid.  Most  of  the  metals 
belong  to  this  class;  mercury,  copper,  silver,  gold,  lead,  thallium,  and  some  others,  set 
free  a less  amount  of  heat,  and  hence  are  not  able  to  decompose  dilute  hydrochloric  acid. 
The  slight  quantity  of  heat  developed  in  the  formation  of  hydrogen  sulphide, 

S,H,  = 4.5, 

indicates  that  the  same  is  readily  decomposed  by  all  the  metals.  In  the  same  way,  by 
adding  the  heat  of  solution, 

S,H2,Aq  = 9.2, 

we  can  easily  ascertain  which  metals  are  precipitated  by  hydrogen  sulphide  from  their 
chlorides,  etc. 

If  in  the  thermo-chemical  equation, 

(Me,Cl„Aq)  (S,H2,Aq)  = (Me,S)  -f  2(II,Cl,Aq), 

the  sum  of  the  heat  developed  upon  the  right  side  is  greater  than  that  upon  the  left,  the 
reaction  will  occur  (precipitation  of  metallic  sulphides)  ; in  the  opposite  case  the  sulphide 
is  decomposed  by  the  dilute  hydrochloric  acid. 


COPPER,  SILVER,  AND  GOLD. 

Considering  the  magnitude  of  their  atomic  weights,  copper,  silver,  and 
gold,  bear  the  same  relation  to  the  alkali  group,  especially  to  sodium,  as 
zinc,  cadmium,  and  mercury  bear  to  magnesium  : 


Na  = 23.015 

Cu=  63.6 
Ag=  107.93 
Au  = 197.2 


Mg=  24.36 
Zn  = 65.4 
Cd  =112 
Hg  = 200.3. 


They  occupy  an  entirely  analogous  position  in  the  three  great  periods  of 
the  periodic  system  of  the  elements  (p.  246),  and  constitute  the  transition 
from  the  elements  of  group  VIII,  especially  from  cobalt,  palladium,  and 
platinum,  to  the  elements  of  group  II — zinc,  cadmium,  and  mercury  : 


Co=  59 
Pd  — 106 
Pt  = 194.8 


Cu  = 63.6 
Ag^  107.93 
Au  = 197.2 


Zn  = 65.4 
Cd  =112 
Ilg  = 200.3. 


This  intermediate  position  of  the  three  elements  about  to  be  discussed 
is  clearly  shown  in  their  entire  physical  deportment.  While  the  ele- 
ments of  group  VIII,  with  the  last  members,  cobalt,  palladium,  and 
])latinum,  fuse  with  difficulty  and  do  not  volatilize,  copper,  silver, 
and  gold  constitute  the  transition  to  the  readily  fusible  and  volatile 
elements,  zinc,  cadmium,  and  mercury.  They  take  an  intermediate 
I)osition,  too,  with  reference  to  their  coefficients  of  expansion,  their 


COPPER,  SILVER,  AND  GOLD. 


329 


atomic  volumes,  and  other  physical  properties.  It  is  noteworthy  that  the 
ability  to  conduct  heat  and  electricity  attains  its  maximum  in  copper,  sil- 
ver, and  gold. 

Not  only  are  the  properties  of  the  free  elements  determined  by  the 
position  of  the  latter  in  the  periodic  system,  but  those  of  their  deriva- 
tives, and  especially  such  as  depend  upon  the  valence  of  the  elements,  are 
influenced  to  a marked  degree  by  the  above  relation.  In  consequence  of 
the  double  periodicity  of  the  great  periods,  copper,  silver,  and  gold  attach 
themselves  to  group  I,  and  especially  to  sodium,  just  as  the  elements  im- 
mediately following,  zinc,  cadmium  and  mercury,  arrange  themselves  with 
group  II  and  magnesium.  Hence  we  find  copper,  silver,  and  gold,  like 

I 

sodium,  yielding  compounds  of  the  form  MeX,  in  which  they  appear  univa- 
lent. Some  of  these  are  isomorphous,  e.g.,  silver  sulphate,  Ag2S04,  with 
sodium  sulphate,  Na2SO^;  sodium  chloride,  NaCl,  cuprous  chloride,  CuCl, 
and  silver  chloride,  AgCl,  crystallize  in  forms  of  the  regular  system. 

But  we  may  say  that  the  similarity  of  copper,  silver,  and  gold  to 
sodium  is  confined  to  these  few  external  properties.  Just  as  the  heavy 
metals,  zinc,  cadmium,  and  mercury  differ  in  many  properties  from  the 
light  metal  magnesium  (p.  313),  so  do  the  metals  copper,  silver,  and 
gold,  possessing  a high  specific  gravity,  distinguish  themselves  in  a still 
higher  degree  from  the  light  metal  sodium.  They  possess  all  the  prop- 
erties belonging  to  the  heavy  metals,  which  are  mainly  characterized  by 
the  insolubility  of  the  oxides,  sulphides,  and  many  salts,  so  that  they 
arrange  themselves  in  a less  marked  degree  with  the  alkali  metals. 

I 

In  the  compounds  constituted  according  to  the  form  MeX,  in  which 
copper,  silver,  and  gold  appear  univalent,  they  exhibit  great  similarity. 
Thus,  the  chlorides,  CuCl,  AgCl,  and  AuCl,  are  white  and  insoluble  in 
water;  soluble,  however,  in  concentrated  hydrochloric  acid,  ammonia, 
the  alkaline  hyposulphites,  etc.,  and  form  double  chlorides,  very  similar 
to  each  other,  with  other  chlorides.  While  silver  only  enters  compounds 

I 

of  the  form  AgX,  copper  and  gold  are  capable  of  yielding  another  form ; 

I II 

copper  forms,  besides  cuprous,  CuX,  also  cupric,  CUX2,  derivatives,  in 
which  it  appears  to  be  bivalent.  The  latter  are  much  more  stable  than  the 
former,  and  embrace  the  ordinary  copper  salts.  Gold,  however,  besides 

I III 

furnishing  aureus,  AuX,  compounds,  has  aurA  derivatives,  AuXg,  in  which 
the  metal  appears  to  be  trivalent  (see  the  Thallium  Group). 

While  copper  and  gold,  in  their  lower  forms,  are  analogous  to  silver  (and 
in  less  degree,  to  sodium),  the  cuprf<:  derivatives  show  a great  resemblance 
to  the  compounds  of  the  metals  of  the  magnesium  group  and  other  metals 
in  their  bivalent  combinations.  Thus,  the  sulphates  of  zinc,  magnesium, 
cupric  oxide  (CuO),  ferrous  oxide  (FeO),  nickelous  oxide  (NiO),  cobalt- 
ous  oxide  (CoO),  and  manganous  oxide  (MnO),  are  similarly  constituted, 
resemble  each  other,  are  isomorphous,  and  form  entirely  analogous  double 
salts  (p.  315)  with  the  alkaline  sulphates.  In  the  same  way  the  carbonates 

II  II 

(MeCOg),  the  chlorates  and  bromates  (MeCl20g  -f  6H2O)  and  others,  are 
28 


330  INORGANIC  CHEMISTRY. 

similarly  conslituted  and  isomorphons.  In  its  aur/c  derivatives,  gold  cx- 

III 

hibits  some  similarity  to  the  aluminium  compounds  (MX  ),  to  those  of 

III 

indium  (InX,,)  and  other  metals,  in  their  trivalent  combinations.  Here 
we  see,  as  already  ob.served  with  mercury  (j).  321),  that  the  similarity  of 
ilie  compounds  of  the  metals  is  influenced  by  the  similarity  of  forms  or  types, 
according  to  which  they  are  composed,  i.  e.,  by  the  valence  of  the  metals. 
If  a metal  form  several  series  of  compounds  of  different  types,  each  series 
is  usually  more  or  less  similar  to  the  compounds  of  other  metals  of  like 
type.  In  this  manner  is  shown  the  resemblance  of  the  compounds  of  the 
following  types ; 


Na^C 

AggO 

Cu.^C) 

Aiqf) 

TljO 

Sodium  oxide. 

Silver  oxide. 

Cuprous  oxide. 

Aurous  oxide. 

Thallous  oxide. 

MgO 

Zn(.) 

CuO 

Fet ) 

IlgO 

Magnesium  oxide. 

Zinc  oxide. 

Cupric  oxide. 

Ferrous  oxide. 

Mercuric  oxide. 

AiPs  Au,03  TI,()3. 

Aluminium  oxide.  I-erric  oxide.  Auric  oxide.  Tliallic  oxide. 


The  character  of  their  derivatives,  varying  with  the  degree  of  combina- 
tion or  valence,  becomes  quite  marked  with  chromium,  manganese  and 
iron,  as  we  shall  later  see.  The  heavy  metals  also  exhibit  a strongly 
positive  and  basic  character  in  their  univalent  combinations.  Thus,  silver 
oxide  (Ag.^O)  and  thallous  oxide  (T^O)  are  strong  bases,  forming  neutral- 
reacting salts  with  acids,  and  even  cuprous  and  aurous  oxides  are  more 
strongly  basic  than  their  higher  forms  of  oxidation.  The  metalloidal 
character  of  the  metals  and  the  acid  nature  of  their  oxides  begin  to  appear 
in  their  trivalent  combinations.  Thus,  in  the  hydroxyl  derivatives  of 
aluminium,  indium,  and  gold,  A1(0H)3,  In(OH)3,  Au(0H)3,  hydrogen 
may  be  replaced  by  the  alkalies  just  as  in  boric  acid,  B(OH)3.  Their 
higher  forms  of  oxidation  show,  like  those  of  the  metalloids,  a pronounced 
acid-like  character  (as  Pb02,  PtOg,  CrOg,  FeOg)  which  is  only  lessened  by 
a high  atomic  weight  of  the  metal  (as  in  PbOg  and  PtOg). 

The  heats  of  formation  of  the  copper,  silver  and  gold  compounds  of  the  MeX  type 
.show  the  same  gradation  as  those  of  the  zinc  group  (p.  327)  : 

(Na,  Cl)  = 97.6  (Ag,  Cl)  = 29.3 

(Cu,  Cl)  = 32.8  (Au,  Cl)  = 5.8. 


COPPER. 

Cu  = 63.6. 

Native  copper  is  found  in  large  quantities  in  America  (Lake  Superior), 
China,  Japan,  also  in  Sweden  and  in  the  Urals.  It  frecpiently  occurs 
crystallized  in  cubes  and  octahedra.  The  most  im|)ortant  and  most  widely 
distributed  of  its  ores  are:  cuprite  (Cu.^O),  malachite  (CuCOg. Cu(OH).2) 


COPPE'K. 


33^ 


and  aziirite  (2CuCO.^.  Cu(OH)2')  (basic  carbonates),  chalcocite  (Cu.^S),  and 
especially  chalco})\  rile  (Ci]2S . Fe2S3)  and  bornite  (3CU2S.  Fe2S3). 

Metallurgy  of  Copper. — 'I'he  extraction  of  copper  from  its  oxygen  ores  is  very  simple  : 
metallic  copper  is  melted  out  when  the  ores  are  ignited  along  with  charcoal.  The  sul- 
phur ores  are  more  difficult  to  work.  i.  The  divided  material  is  first  roasted  in  the  air, 
whereby  the  sulphur  content  is  reduced  : 

3(Cu2S.  Fe2S3)  -f  21O  = 6CuO  -f-  5FeS  -j-  FeO  -f  ySOg. 

2.  The  mass  is  afterward  fused,  when  we  get  the  so-called  coarse  metal,  which  con- 
tains iron  sulphide  in  addition  to  copper  sulphide  : 

6CuO  + sFeS  ^ (3CU2S  + FeS)  + 4FeO  + SO2. 

3.  Crude  copper  (black  copper)  is  obtained  by  further  roasting  and  by  the  “fusion 
reaction”  of  this  mass: 

3CU2S  + FeS  -f  9O  :=  3CU2O  -f  FeS  -f  3SO2  = 3CU2  + FeO  -f  4SO2. 

The  ferrous  oxide  formed  here  is  taken  up  by  the  silica  to  form  a slag.  In  the  more 
recently  constructed  works,  particularly  in  the  United  States,  the  sulphide  ores  are  first 
roasted  and  then  melted  in  cupola  or  reverberatory  furnaces,  after  which  the  coarse  metal 
is  subjected  to  the  Bessemer  process  and  the  resulting  black  copper  is  electrolyzed.  See 
Chem.  Zeit.  1895,  977  ; 1897,  995  ; also  Jahrb.  d.  Chem.  v (1895',  287. 

In  refining  copper  electrcilytically,  plates  cast  from  black  copper  are  suspended  as 
anodes  in  the  solution  of  copper  sulphate.  The  kathodes  are  plates  of  pure  copper.  The 
copper  is  dissolved  from  off  the  anode  and  deposited  on  the  kathode.  Impurities,  includ- 
ing any  of  the  nobler  metals,  settle  at  the  bottom  as  mud. 

Copper  can  also  be  obtained  in  the  wet  way  from  its  ores.  The  pulverized  ore  is 
treated  with  solutions  of  ferric  chloride  or  sulphate,  when  the  copper  sulphide  is  trans- 
posed to  soluble  copper  salts  in  the  sense  of  the  following  equations  : 

CU2S  -f  Fe2Clg  = 2FeCl2  + CU2CI2  -j-  S ; 

CuS  -f  Fe2Clg  = CuCb  f-  2FeCl2  -}-  S ; 

CU2S  4-  2Fe2(SOb3  = 2CUSO4  -f  4FeS04  S. 

The  copper  is  then  precipitated  electrolytically  from  the  solution. 

The  following  countries  are  rich  in  copper  : United  States  (Lake  Superior  and  Mon- 
tana), Spain  and  Portugal,  also  Chili,  Japan,  and  Germany  (Prussia,  Mansfield). 

The  following  are  excellent  reference  works  on  the  metallurgy  of  copper  : Ost’s  Tech- 
nologie,  Dammer’s  Technologie  (p.  190),  and  Durre’s  Vorlesungen  fiber  allgemeine 
Hfittenkunde  (Halle,  1898). 

To  obtain  chemically  pure  copper,  the  pure  oxide  is  heated  in  a stream 
of  hydrogen,  or  the  solution  of  copper  sulphate  is  decomposed  by 
electrolysis. 

Metallic  copper  possesses  a characteristic  red  color,  and  transmits  a 
green  light  in  thin  leaflets.  It  is  rather  soft  and  ductile,  and  possesses  a 
specific  gravity  of  8. 5-8. 9.  It  melts  at  about  1080°,  and  vaporizes  in 
the  oxyhydrogen  flame.  It  remains  unaltered  in  dry  air;  in  moist,  it  is 
gradually  coated  with  a green  layer  of  basic  copper  carbonate.  When 
heated  in  the  air,  it  oxidizes  to  black  cupric  oxide. 

Copper  is  not  changed  by  dilute  hydrochloric  or  sulphuric  acids;  if  it 
be  moistened  with  these,  and  exj)osed  to  the  air,  it  absorbs  oxygen,  and 
gradually  dissolves.  It  issimilarly  dissolved  by  ammonium  hydroxide.  Hot 
concentrated  sulphuric  acid  converts  it  into  copper  sulphate,  with  evolu- 


33^ 


INORGANIC  CHEMISTRY. 


tion  of  sulphur  dioxide.  It  dissolves  in  dilute  nitric  acid  in  the  cold, 
with  evolution  of  nitric  oxide.  Zinc,  iron  and  al.so  jihosphorus  precipi- 
tate metallic  copper  from  the  aqueous  solutions  of  its  salts. 


Copper  forms  two  series  of  compounds,  known  as  cu\nous  and  cupr/V. 
In  the  latter  co})})er  is  bivalent : 

CuO  CuGj  Cu(OII),  CuSO,. 

These  are  more  stable  than  the  cuprous  derivatives;  the  ordinary  copper 
salts  belong  to  them.  In  many  respects  they  resemble  the  compounds 
of  other  dyad  metals,  especially  those  of  the  magnesium  group,  and  the 
ous  comj)ounds  of  iron  (FeO),  manganese  (MnO),  cobalt  and  nickel 
(see  p.  329). 

The  cuj)rous  compounds  are,  on  the  other  hand,  very  unstable,  absorb 
oxygen  from  the  air,  and  pass  into  cupric  derivatives.  They  show  some 
similarity  to  the  mercurous  derivatives  (p.  323),  and  possess  an  analogous 
composition  : 

CuCl  Cul  CU2O  Cu.^S. 

Oxygen  salts  of  cuprous  oxide  are  not  known. 

From  the  formulas  given  above,  copper,  like  silver,  is  univalent  in  the  cuprous  com- 
pounds. This  is  also  experimentally  indicated  by  the  boiling  and  melting  points  shown 
by  solutions  of  cuprous  salts  in  pyridine,  piperidine,  and  methyl  sulphide.  The  salts 
are  not  dissociated  in  these  solutions  ; they  are  non-conductors.  From  what  has  been 
said  on  p.  268  some  of  them  probably  have  the  simple  formula,  especially  cuprous 
bromide,  CuBr. 

Other  salts  are  disposed  to  appear  as  double  molecules  ; the  cyanide  is  only  known  as 
Cu2(CN)2'.  We  can  therefore  assume,  as  was  done  in  the  case  of  the  mercurous  salts, 
the  presence  of  a bivalent  group  of  two  metallic  atoms  : 

CN-Cu-Cu-CN  Cl-Cu-Cu-Cl. 

The  vapor  density  of  cuprous  chloride  even  at  1 700°  corresponds  to  the  formula  CU2CI2. 
Both  formulas  should  be  given  equal  value.  [See  pp.  322,  324,  339,  and  Werner,  Z.  f. 
anorg.  Ch.  15  (1897)  l]. 


CUPROUS  COMPOUNDS. 

Cuprous  Oxide,  CiqO,  occurs  as  cuprite  crystallized  in  regular 
octahedra.  It  is  obtained  artificially  by  boiling  a solution  of  copper  sul- 
phate and  grape  sugar  with  potassium  hydroxide,  when  it  separates  as  a 
crystalline,  bright  red  powder.  It  does  not  change  in  the  air,  and  is 
readily  soluble  in  ammonium  hydroxide.  The  solution  absorbs  oxygen, 
and  while  forming  cupric  oxide  acquires  a blue  color.  By  the  action  of 
sulphuric  and  other  oxygen  acids,  it  forms  cupric  salts,  half  of  the  copper 
sejjarating  as  metal : 

CU2O  -f  II2SO,  CuSO,  + Cu  + H2O. 

d'he  hydroxide^  Cu.2(OH)2,  is  precipitated  by  the  alkalies  as  a yellow  pow- 


CUPROUS  COMPOUNDS.  333 

der  from  hydrochloric  acid  solutions  of  cuprous  chloride,  CU2CI2.  It 
oxidizes  in  the  air  to  cupric  hydroxide. 

Cuprous  Chloride,  CuC.l  or  CU2CI2,  is  produced  by  the  combustion 
of  metallic  copper  in  chlorine  gas  (together  with  cupric  chloride,  CuClj), 
upon  conducting  hydrochloric  acid  over  copper  at  alow,  red  heat;  by 
boiling  the  solution  of  cupric  chloride  with  copper: 


CuCb  -|-  Cu  = CU2CI2, 

and  by  the  action  of  many  reducing  substances  upon  cupric  chloride.  It 
is  most  conveniently  made  by  passing  sulphur  dioxide  through  a concen- 
trated solution  of  copper  sulphate  and  sodium  chloride,  when  it  separates 
as  a white,  shining  powder,  consisting  of  small  tetrahedra.  It  fuses  at 
430°,  and  distils  at  about  1000°;  its  vapor  density  at  1600-1700°  corre- 
sponds to  the  formula  CujClj.  In  moist  air  it  rapidly  becomes  green* 
owing  to  oxygen  absorption,  and  the  formation  of  basic  cupric  chloride. 
Cl 

Cu<oH-  becomes  black  on  exposure  to  light.  Cuprous  chloride  is 

readily  soluble  in  concentrated  hydrochloric  acid  and  in  ammonium 
hydroxide;  both  solutions  jjossess  the  characteristic  property  of  absorb- 
ing carbon  monoxide.  Colorless  leaflets  separate  from  solutions  saturated 
with  the  gas.  Their  composition  seems  to  be  CuaCb  -f  CO  -j-  2H2O. 

Cuprous  Iodide,  Cul  or  Ci]2l2,  is  precipitated  from  soluble  cupric 
salts  by  potassium  iodide  : 

2CuS0^  -f  4KI  = CU2I2  + 2X280^  + I2. 


By  extracting  the  co-precipitated  iodine  by  means  of  ether  it  is 
obtained  as  a gray  powder,  insoluble  in  acids. 

Cuprous  Sulphide,  CU2S,  occurs  as  chalcocite  crystallized  in  forms  of 
the  rhombic  system.  It  is  produced  by  burning  copper  in  sulphur  vapor, 
and  by  heating  cupric  sulphide  in  a current  of  hydrogen  ; after  fusion  it 
solidifies  in  crystals  of  the  regular  system.  Combined  with  silver  sul- 

Cu  'I 

phide  it  constitutes  the  mineral  stromeyerite,  | S or  CU2S.  Ag2S,  iso- 
morphouswith  chalcocite. 

Copper  Hydride,  CuH  or  CU2H2,  belongs  to  the  derivatives  of 
univalent  copper.  Jf  a solution  of  copper  sulphate  be  digested  with 
hypophosphorous  acid,  the  hydride  separates  as  a yellow  amorphous  pre- 
cipitate which  soon  acquires  a brown  color.  At  60°  it  decomposes  into 
copper  and  hydrogen. 

With  hydrochloric  acid  it  forms  cuprous  chloride  : 


CU2H2  + 2HCI  = Cu2a2  + 2H2. 


(Compare  Mylius  and  Fromm,  Ber.  27  (1894),  i,  647.) 


334 


INORGANIC  CHKMIS'IRV. 


CUPRIC  COMPOUNDS. 

'riie  cuiiric  salts,  when  liydrous,  are  generally  colored  blue  or  green  ; 
they  are  colorless  in  the  anhydrous  condition. 

Cupric  Hydroxide,  Cu(()ll),^,  separates  as  a voluminous  bluish  pre- 
cipitate when  sodium  or  potassium  hydroxide  is  added  to  soluble  copper 
salts.  \Vhen  heated,  even  under  water,  and  particularly  in  the  jiresence 
of  alkali,  it  loses  water,  and  is  changed  to  black  cupric  oxide.  In  the 
presence  of  alkaline  chlorides,  or  of  cui)ric  chloride  it  is  i)artly  reduced 
to  a cuprous  salt  with  the  production  of  an  alkaline  peroxide  in  the 
evolution  of  oxygen  : 

2Cu(0II),  + 2KBr  = CujBr,  + K./)^  f 2ll,0. 

(Spring  and  Liicion,  Z.  f.  anorg.  Ch.  2 (1892),  195). 

Cupric  Oxide,  CuO,  is  usually  obtained  by  the  ignition  of  copper 
turnings  in  the  air,  or  by  heating  cupric  nitrate.  It  forms  a black 
amorphous  powder,  which,  at  higher  temperatures,  settles  together  and 
acquires  a metallic  luster.  By  heating  with  organic  substances  their 
carbon  is  converted  into  carbon  dioxide,  and  the  hydrogen  into  water, 
th-;  ciq^ric  oxide  being  reduced  to  metal ; upon  this  rests  the  application 
of  cupric  oxide  in  the  analysis  of  such  comj)ounds. 

Copper  oxide  and  hydroxide  dissolve  in  ammonium  hydroxide  with  a 
dark-blue  color.  The  solution  possesses  the  power  of  dissolving  wood 
fiber  (cotton-wool,  linen,  filter-paper,  etc.) — Schweizer' s reagent. 

Cupric  Chloride,  CUCI2,  is  formed  by  the  solution  of  cupric  oxide 
or  carbonate  in  hydrochloric  acid.  It  crystallizes  from  aqueous  solution, 
with  two  molecules  of  water,  in  bright  green  rhombic  needles,  and  is 
readily  soluble  in  water  and  alcohol.  When  heated,  it  parts  with  its 
water,  becoming  anhydrous  chloride,  which  at  a red  heat  is  decomposed 
into  chlorine  and  cuprous  chloride.  It  yields  beautifully  crystallized 
double  salts  with  potassium  and  ammonium  chlorides.  Cupric  bromide 
is  like  the  chloride ; the  iodide  is  not  known,  since  in  its  formation  it  at 
once  breaks  down  into  cuprous  iodide  and  iodine. 

Copper  Sulphate,  CuSO^  + cupric  suli)hate,  copper  vitriol — 

may  be  obtained  by  the  solution  of  copper  in  concentrated  sulphuric 
acid.  It  is  produced  on  a large  scale  by  roasting  chalcocite.  It  forms 
large  blue  crystals  of  the  triclinic  system,  which  effloresce  somewhat  upon 
exposure.  At  100°  the  salt  loses  four  molecules  of  water  ; the  fifth  sepa- 
rates above  200°.  The  anhydrous  sulphate  is  colorless,  absorbs  water 
very  energetically,  and  returns  to  the  blue  hydrous  compound. 

Although  copi)er  sulphate  only  crystallizes  with  five  molecules  of  water, 
it  is  capable,  like  the  sulphates  of  the  magnesium  group,  of  forming 
double  salts  with  potassium  and  ammonium  sulphates,  which  crystallize 
with  six  molecules  of  water,  and  are  isomorphous  with  the  double  salts 
of  the  metals  of  the  magnesium  group. 

('oppcr  sulphate  is  em[)loyed  in  electro-plating.  When  its  solution  is 
decomposed  by  the  galvanic  current  copjier  separates  at  the  negative 
])ole,  and  deposits  in  a regular  layer  upon  the  conducting  objects  con- 


CUPRIC  COMPOUNDS. 


335 


nected  wiih  the  electrode.  The  positive  pole  is  a copperplate:  the 
ions,  SO^,  dissolve  a quantity  of  it  equal  to  the  copper  deposited  on  the 
negative  pole;  hence  the  solution  always  has  the  same  concentration. 

Ammonium  hydroxide  added  to  a copper  sulphate  solution  in  sufficient 
quantity  to  dissolve  the  cupric  hydroxide  produced  at  first,  changes  the 
color  of  the  liquid  to  a dark  blue.  From  this  solution  alcohol  precipitates 
a dark-blue  crystalline  mass  with  the  composition  CuSO^.qNHg  -}-  H2O. 
Heated  to  150°  this  compound  loses  water  and  two  molecules  of  ammonia, 
and  becomes  CUSO4.2NH3.  It  is  supposed  that  these  compounds  are 
ammonium  salts  in  which  a part  of  the  hydrogen  is  replaced  by  copper; 
they  have  been  designated  cuprammoniimi  compounds,  e.  g.  : 

Cuprammonium  sulphate. 

The  other  soluble  copper  salts  yield  similar  compounds  with  ammonium 
hydroxide. 

Cupric  Nitrate,  Cu(N03)2,  crystallizes  with  three  or  six  molecules  of 
water,  has  a dark-blue  color  and  is  readily  soluble  in  water  and  alcohol. 
Heat  converts  it  into  cupric  oxide. 

Copper  Carbonates. — The  neutral  salt  (C11CO3)  is  not  known. 
When  sodium  carbonate  is  added  to  a warm  solution  of  a copper  salt  the 

basic  carbonate,  CuCOg . Cu(OH)2,  or  separates  as  a 

green  precipitate.  It  occurs  in  nature  as  malachite,  which  is  especially 
abundant  in  Siberia.  Another  basic  salt,  aCOgCu . Cu(OH  ^2,  is  the  beau- 
tiful blue  nzurite. 

Copper  Arsenite  separates  as  a beautiful  bright  green  precipitate, 
upon  the  addition  of  sodium  arsenite  to  a copper  solution.  This  preci- 
])itate  has  never  been  obtained  in  homogeneous  form  and  of  a definite 
composition.  It  was  formerly  employed  as  a pigment,  under  the  name 
of  Scheele' s green,  but  at  present,  owing  to  its  poisonous  character,  it  has 
been  replaced  by  other  green  colors  (Guignet’s  green  and  aniline  green). 

Cupric  Sulphide,  CuS,  is  a black  compound,  precipitated  from  cop- 
per solutions  by  hydrogen  sulphide.  It  is  insoluble  in  dilute  acids. 
When  moist,  it  slowly  oxidizes  in  the  air  to  cupric  sulphate.  Heated  in 
a stream  of  hydrogen,  it  forms  cuprous  sulphide,  CUgS. 


Alloys  of  Copper. — Pure  copper  is  very  ductile,  and  may  be  readily 
rolled,  and  drawn  out  into  a fine  wire.  It  does  not  make  good  cast- 
ings, because  it  contracts  unequally  upon  cooling  and  does  not  fill  out 
the  moulds.  For  such  pur])oses,  alloys  of  copper  are  em]:)loyed  which, 
in  addition,  possess  other  technically  valuable  properties.  The  most 
important  copper  alloys  are: 

Brass,  consisting  of  three  parts  of  copjjer  and  one  part  of  zinc.  It  has 
a yellow  color,  and  is  considerably  harder  than  ]jure  co])per.  Ordinarily, 
one  or  two  per  cent,  of  lead  are  added  to  the  brass,  which  facilitates  its 


33® 


INORGANIC  CHEMISTRY. 


working  upon  the  turning-lathe.  Tombac  contains  15  per  cent,  of  zinc, 
and  has  a gold-like  color.  The  alloy  of  i part  of  zinc  and  5.5  parts  of 
copper  answers  for  the  manufacture  of  spurious  gold  leaf,  d'he  alloys  of 
copper  with  tin  are  called  bronzes.  Most  of  the  modern  bronzes  also 
contain  zinc  and  lead;  those  from  Japan,  gold  and  silver.  The  cannon 
bronze  contains  90  per  cent,  of  coi)per  and  10  per  cent,  of  tin  ; bell-metal 
has  20-25  cent,  of  tin. 

Argentan  is  an  alloy  of  copper,  zinc,  and  nickel  (see  Nickel).  The 
German  cop])er  coins  consist  of  95  per  cent,  of  coj)jjer,  4 [)er  cent,  of  tin, 
and  I per  cent,  of  zinc. 

Of  the  more  recently  introduced  alloys  of  copi)er  we  may  mention  : 

Phosphor-bi'Ofize. — This  consists  of  90  parts  of  cop|)er,  9 parts  of  tin 
and  0.5-0. 8 parts  of  phosphorus.  By  the  last  ingredient  the  bronze  is 
increased  in  hardness,  and  its  solidity  and  resistance  to  oxidation  are  also 
increased.  It  is  em})loyed  in  making  axle  bearings  and  various  parts  of 
machinery. 

Silicon  bronze,  containing  silicon  instead  of  phosphorus,  is  charac- 
terized by  great  firmness  and  conductivity.  It  is  used  for  telephone 
wires  and  as  conductii'g  cables  for  electric  railways. 

Manganese  bronze  contains  70  per  cent,  of  copi)er  and  30  per  cent,  of 
manganese  (it  is  the  cupro-manganese  of  Letrange),  and  may  be  melted 
with  copper  and  copper  alloys,  im]3arting  to  these  solidity  and  great 
hardness.  It  is  especially  employed  in  making  ship-propellers. 

See  page  347  for  Aluminium  Bronze. 


Recognition  of  Copper  Compounds. — Most  copper  compounds 
containing  water  have  a blue  or  green  color.  With  the  exception  of 
copper  sulphide  they  all  dissolve  in  ammonium  hydroxide,  with  a blue 
color.  Potassium  ferrocyanide  produces  a reddish-brown  precipitate  or 
color  in  aqueous  solutions  of  cupric  salts.  When  a clean  piece  of  iron  is 
introduced  into  a co|)per  solution,  it  becomes  covered  with  a red  layer 
of  metallic  copper.  Volatile  copper  compounds  tinge  the  flame  blue  or 
green.  The  spectrum  of  such  a flame  is  characterized  by  several  blue  or 
green  lines. 


SILVER. 

Ag  = 107.93. 

Silver  occurs  native.  Its  most  important  ores  are  argentite  or  silver 
glance  (silver  sulphide,  Ag^S)  and  various  compounds  with  sulphur, 
arsenic,  antimony,  copper  and  other  metals  (light-red  silver  ore,  prous- 
tite,  Ag^AsSj,  and  dark-red  silver  ore,  pyrargyrite,  Ag3SbS3).  Of  rarer 
occurrence  are  combinations  with  chlorine  (horn-silver,  AgCl),  bromine 
and  iodine.  Slight  quantities  of  silver  sulphide  are  ])resent  in  almost 
every  galenite  (PbS).  The  principal  localities  for  silver  ores  are  the 
United  States,  Mexico,  Bolivia,  Australia,  Saxony  and  Chili. 


SILVER. 


337 


Metalliirf:;y  of  Silver. — The  metallurgy  of  silver  ores  involves  methods  dependent  upon 
the  silver  content  of  the  ore,  the  chemical  constitution  of  the  latter  and  the  conditions  of 
locality.  It  is  usually  a complicated  procedure.  The  important  methods  alone  will  be 
indicated  briefly  here  ; details  must  be  sought  in  special  works  upon  chemical  technology 
(p.  331).  It  is  only  the  richest  ores  from  which  the  silver  can  be  directly  recovered  ; 
most  of  the  metal  is  obtained  either  (i)  by  first  preparing  an  argentiferous  lead  and  then 
separating  these  two  metals  subsequently  ; or  {2)  by  the  amalgamation  process  ; or  (3)  by 
converting  the  silver  into  soluble  compounds  and  precipitating  it  from  the  solutions. 

1.  The  ores  are  worked  together  with  lead  ores  for  metal  or,  if  they  contain  more  than 
yL  of  silver,  they  are  fused  directly  with  lead.  Various  methods  are  pursued  to  enrich 
the  silver-lead  alloy  with  the  first  metal,  {a)  Pattinsouiziug.  When  lead  containing 
silver  is  fused  and  cooled  slowly  at  first  pure  lead  crystallizes  out  which  is  removed  with 
sieves;  an  alloy  rich  in  silver  and  readily  fused  remains.  {1))  Parkes's  process.  This 
method  consi.sts  in  adding  zinc  to  the  silver-lead  alloy  of  low  silver  content  and  obtaining 
thereby  an  alloy  of  lead-zinc-silver,  fusible  with  difficulty.  When  the  fusion  cools  this 
alloy  appears  on  the  surface  as  “zinc  crust.”  The  zinc  is  expelled  from  it  by  distillation. 
The  lead  now  rich  in  silver  (obtained  by  one  or  the  other  method)  is  cupelled,  i.  e.,  it  is 
heated  with  air-access  in  a furnace  the  bottom  of  which  consists  of  a porous  mass.  The 
lead  is  oxidized  to  lead  oxide  which  partly  escapes  and  is  in  part  absorbed.  The  silver 
is  not  oxidized  and  remains  in  the  cupel. 

2.  The  amalgamation  process  is  based  on  the  fact  that  mercury  dissolves  silver  and 
also  decomposes  its  compounds  with  the  formation  of  an  amalgam. 

The  ore  is  roasted  with  sodium  chloride,  whereby  silver  chloride  is  produced.  The 
divided  material  is  then  mixed  with  iron  scraps  and  water  in  rotating  vessels.  The  iron 
causes  the  precipitation  of  the  metallic  silver  from  its  chloride  : 

2AgCl  -|-  Fe  = FeCl2  -f  2Ag. 

To  free  the  metal  from  various  impurities  it  is  dissolved  in  mercury  and  the  liquid 
amalgam  ignited,  when  mercury  distils  off  and  silver  remains.  Owing  to  scarcity  of  com- 
bustible material,  the  conversion  of  silver  ores  into  silver  chloride  is  executed,  in  Mexico 
and  Peru,  by  mixing  the  ores  with  sodium  chloride  and  copper  sulphate  in  the  presence 
of  water.  In  this  way  cuprous  chloride  is  produced,  which  is  transposed,  with  silver  sul- 
phide, into  silver  chloride  and  cuprous  sulphide  : 

CU2CI2  + Ag2S  = CU2S  + 2AgCl. 

3.  By  solution  and  precipitation.  The  ores  are  roasted  alone  and  afterwards  together 
with  salt.  Silver  sulphate  is  first  produced  and  then  silver  chloride.  The  roasted  mass 
is  extracted  with  a concentrated  salt  solution  which  dissolves  the  silver  chloride,  and 
from  this  solution  the  silver  metal  is  precipitated  by  metallic  copper.  The  resulting 
cuprous  chloride  solution  is  acted  upon  with  iron  to  recover  the  copper  (Augustin).  Or 
after  roasting  the  ore  alone  silver  sulphate  is  dissolved  out  with  water  and  the  solution 
treated  with  copper  (Ziervogel).  The  silver  can  also  be  dissolved  out  by  sodium  copper 
thio.sulphate,  the  metal  being  then  precipitated  by  means  of  a calcium  sulphydrate  solu- 
tion (Russell). 

Commercial  silver  (work-silver)  is  not  pure,  but  invariably  contains 
copper  and  traces  of  other  metals  in  greater  or  less  quantity.  To  pre- 
pare chemically  pure  metal,  the  work-silver  is  dissolved  in  nitric  acid, 
and  from  the  solution  of  nitrates  thus  obtained  hydrochloric  acid  pre- 
cipitates the  silver  as  chloride  : 

AgNOg  -f  HCl  = AgCl  + HNO3. 

The  latter  is  reduced  by  various  methods;  either  by  fusion  with  sodium 
carbonate,  or  by  the  action  of  zinc  or  iron  in  the  presence  of  water: 

2AgCl  -f  Na-^CO,  2NaCl  -f-  2Ag  + COg  | O. 

2AgCl  -f-  Zn  = ZnClj  T 2Ag. 


29 


338 


INORGANIC  CHEMISTRY. 


Silver  is  a pure  white,  brilliant  metal,  of  specific  gravity  10.5.  It  is 
tolerably  soft  and  very  ductile,  and  can  be  drawn  out  to  a fine  wire.  It 
crystallizes  in  regular  octahedra.  It  fuses  at  960°,  and  is  converted  into 
a greenish  vapor  in  the  oxyhydrogen  flame.  It  is  the  best  conductor  of 
heat  and  electricity.  Silver  is  not  oxidized  by  oxygen  ; by  the  action  of 
ozone  it  is  covered  with  a very  thin  layer  of  silver  peroxide.  When  in 
molten  condition,  silver  absorbs  22  volumes  of  oxygen  without  combin- 
ing chemically  with  it ; the  absorbed  gas  escapes  again  when  the  metal 
cools  [sprouting  or  si)itting  of  silver]. 

Silver  is  capable  of  existing  in  three  allotropic  forms,  which  have 
properties  greatly  different  from  those  of  ordinary  silver.  The  first  form 
is  soluble  in  water  and  has  a blue  color.  The  second  variety  is  insoluble, 
and  somewhat  resembles  the  first  form.  The  third  closely  resembles  gold 
in  color  and  luster.  These  allotropic  varieties  of  silver  are  broadly  dis- 
tinguished from  normal  silver  by  color.  They  very  likely  are  more  active 
conditions  of  silver,  common  silver  being  a polymerized  variety  (Am.  Jr. 
Science,  37,  476). 

Silver  unites  directly  with  the  halogens;  by  the  action  of  hydrochloric 
acid  it  becomes  coated  with  an  insoluble  layer  of  silver  chloride.  Boiled 
with  strong  sul])huric  acid,  it  dissolves  to  sulphate: 

^■^8  T zIIjSO^  = Ag.^SO^  4“  SO2  T 2II2O. 

The  best  solvent  of  silver  is  nitric  acid,  which  even  in  a dilute  state 
and  unaided  by  heat,  converts  it  into  nitrate. 

As  silver  is  rather  soft,  it  is  usually  employed  in  the  arts  alloyed  with 
copper,  whereby  it  acquires  a greater  hardness.  Most  silver  coins  consist 
of  90  per  cent,  of  silver  and  10  per  cent,  of  copper;  the  English  shil- 
lings contain  92.5  per  cent,  of  silver.  What  is  termed  the  fineness  of 
silver  is  the  number  of  parts  per  thousand  which  the  alloy  contains. 


Oxygen  forms  three  compounds  with  silver,  but  only  the  oxide  yields 
corresponding  salts. 

Silver  Oxide,  AgjO,  is  thrown  out  of  a silver  nitrate  solution  by 
sodium  or  potassium  hydroxide  as  a dark-brown,  amorphous  precipitate. 
It  is  somewhat  soluble  in  water,  and  blues  red  litmus-paper.  In  this,  and 
in  the  neutral  reaction  of  the  nitrate,  the  strongly  basic,  alkaline  nature 
of  silver  and  its  oxide  exhibits  itself;  the  soluble  salts  of  nearly  all  of 
tlie  other  heavy  metals  show  an  acid  reaction  (p.  330).  When  heated  to 
250°,  the  oxide  decomposes  into  metal  and  oxygen  ; at  too°  it  is  reduced 
by  hydrogen.  Silver  hydroxide,  AgOH,  is  not  known  ; the  moist  oxide 
reacts,  however,  very  much  like  an  hydroxide. 

On  dissolving  precipitated  silver  oxide  in  ammonium  hydroxide,  black 
crystals  (Ag.^O.  2NH3)  separate  when  the  solution  evaix)rates,  and  when 
dry  these  exi)lode  ui)on  the  slightest  disturbance — fulminating  silver. 

Silver  Suboxide,  Ag/),  is  said  to  l)e  produced  by  heating  silver  citrate  in  a current 
of  hydrogen,  and  is  a black,  very  unstable  powder,  which  decomposes  readily  into  silver 
oxide  and  silver. 


SILVER. 


339 


Silver  Peroxide,  AgO  or  Ag202,  is  formed  by  passing  ozone  over  silver  or  its  oxide, 
or  by  the  decomposition  of  the  nitrate  by  the  electric  current.  It  consists  of  black, 
shining  octahedra,  and  at  ioo°  decomposes  into  silver  oxide  and  oxygen  (Mulder). 


The  salts  of  silver  correspond  to  the  oxide  Ag.^O,  and  are  all  constituted  according  to 
the  form  AgX,  hence  are  termed  argentic.  They  are  analogous  to  the  cuprous  and  mer- 
curous derivatives,  and  show  a great  resemblance  to  the  former.  It  would,  therefore,  be 
more  correct  to  designate  them  argent£>«.y.  Compounds  of  the  bivalent  form  AgXj  are 
not  known  for  silver.  If,  however,  the  mercurous  and  cuprous  compounds  are  expressed 
by  double  formulas  (pp.  322,  332)  : 

CuCl  Cu,  llgCl  Hg, 

1 I )0  and  I I )0 

CuCl  Cu-^  HgCl  Hg/ 


which  view  is  supported  by  their  chemical  deportment,  and  is  experimentally  confirmed 
in  the  case  of  cuprous  chloride  by  its  vapor  density  ; those  of  silver  might  be  represented 
by  analogous  formulas : 


AgCl 

I 

AgCl 


Ag\ 

I )o 


AgNOa 

AgN03. 


Then  the  silver  atom  would  be  bivalent  and  a complete  parallelism  would  be  estab- 
lished with  copper.  In  support  of  this  view  there  exists  the  fact  that  when  the  silver 
halides  are  dissolved  in  pyridine,  methyl  sulphide,  etc.,  the  colligative  properties  of 
their  solvents  are  influenced  as  if  halide  molecules  possessing  the  formulas  Ag2Cl2,  Ag2- 
Br2,  Ag2l2,  and  possibly  in  part  higher  formulas,  were  present.  See  pp.  268,  332.  The 
chemical  formulas  of  solid  bodies  do  not  generally  designate  their  true  molecular  values 
as  in  the  case  of  gases,  but  only  their  simplest  atomic  composition.  It  is  very  probable 
that  even  the  simplest  chemical  compounds,  e.  g.^  potassium  chloride  and  silver  chloride, 
consist  in  their  solid  condition  of  complex  molecules  corresponding  to  the  formulas 
(KCl)n,  (AgCl)ni.  An  argument  supporting  this  view  is  afforded  by  the  existence  of 
different  modifications  of  chloride  and  bromide  of  silver ; these  differ  from  each  other  in 
their  external  properties,  and  their  varying  susceptibilities  to  light.  The  doubling  of 
formulas,  as  .shown  above  with  CU2CI2,  Hg2Cl2,  etc.,  is  mainly  due  to  the  tendency  to 
deduce  all  the  compounds  of  an  element  from  a constant  value,  according  to  the  doctrine 
of  con.stant  valence.  This  is,  however,  impossible  (p.  ifp,  322).  According  to  present 
notions  of  valence,  and  as  it  is  presented  in  the  periodic  sy.stem,  compounds  (MeCl, 
MeCl2,  MeCl.j,  etc. ) are  constituted  according  to  definite  forms  or  types  that  may  materi- 
ally determine  their  properties  (p.  322).  So  far  as  the  similarity  of  metallic  compounds 
is  concerned,  it  is  of  .secondary  importance  whether  the  quantities  corresponding  to  the 
simple  formulas,  in  the  solid  or  gaseous  state,  do  unite  to  larger,  complex  molecules 
(compare  HgCl  and  CU2CI2, — AICI3,  A^CHg).^  and  AL^Clf;,  GaClj  and  Ga2Cl^,, — SnCl2, 
80201^,  PbCl2,  etc. ).  In  case  of  the  se.squioxides  Me2G3  it  is  also  immaterial  whether 
they  are  derived  from  supposed  trivalent  elements,  or  from  those  that  are  quadrivalent. 
The  .same  may  be  remarked  of  the  metallic  compounds  Me30^  ^ { MeO.  0)2Me  (see 
Spinels,  351). 

The  use  of  simple  or  of  double  formulas  for  the  metallic  compounds  is  therefore  of  no 
special  importance. 


Silver  Chloride,  AgCl,  exists  in  nature  as  horn -silver.  When  hydro- 
chloric acid  is  added  to  solutions  of  silver  salts,  a white,  curdy  precipitate 
separates  ; the  same  fuses  at  490°  to  a yellow  liquid,  which  solidifies  to  a 
horn-like  mass.  The  chloride  is  insoluble  in  dilute  acids;  it  dissolves 
somewhat  in  concentrated  hydrochloric  acid  and  in  sodium  chloride, 


340 


INORGANIC  CHEMISTRY. 


very  easily  in  ammonium  hydroxide,  i)otassium  cyanide,  and  sodium 
hyposulphite.  It  crystallizes  from  ammoniacal  solutions  in  large,  regular 
octahedra.  Dry  silver  chloride  absorbs  ammonia  gas,  forming  a white 
compound  of  the  formula  2AgC1.3NH3,  which  at  38°  gives  up  its 
ammonia. 

Silver  Bromide,  AgBr,  is  precipitated  from  silver  salts  by  hydro- 
bromic  acid  or  soluble  bromides.  It  has  a bright  yellow  color  and  dis- 
solves with  more  difficulty  than  the  chloride  in  ammonium  hydroxide; 
in  other  respects  it  is  perfectly  similar  to  the  latter.  Heated  in  chlorine 
gas  it  is  converted  into  silver  chloride. 

Silver  Iodide,  Agl,  is  distinguished  from  the  chloride  and  bromide 
by  its  yellow  color,  and  its  insolubility  in  ammonia,  bused  silver  iodide 
at  first  solidifies  in  isometric  crystals,  which  gradually  change  to  hexag- 
onal forms,  but  when  the  latter  are  heated  to  146°,  they  suddenly  revert  to 
the  isometric  forms.  It  dissolves  readily  in  hydriodic  acid,  to  Agl. HI, 
which,  upon  evaporation  of  the  solution,  separates  in  shining  scales. 
Heated  in  chlorine  or  bromine  gas,  it  is  converted  into  chloride  or 
bromide;  conversely,  chloride  and  bromide  of  silver  are  converted  into 
silver  iodide  by  the  action  of  hydriodic  acid. 

These  opposite  reactions  are  explained  by  the  principle  of  the  greatest  evolution  of 
heat.  Chlorine  and  bromine  expel  iodine  from  all  iodides  because  the  heat  of  forma- 
tion of  the  latter  is  less  than  that  of  the  bromides  and  chlorides.  Again,  hydriodic 
acid  (gaseous  or  in  aqueous  solution)  converts  silver  chloride  into  the  iodide  according  to 
the  equation  : 

AgCl  4-  HI  = Agl  -f-  IlCI, 

because  the  heat  modulus  of  the  reaction  is  positive  (for  gaseous  halogen  hydrides 
T12.5  Cal.,  for  the  .solution  -(-10.6  Cal.  See  the  Table  at  close  of  book).  Yet  it  may 
be  noted  here  that  Julius  found  that  silver  chloride  and  iodide  were  changed  to  bromide 
upon  heating  them  in  bromine  vapor,  and  further  that  the  chloride  and  bromide  were 
also  changed  by  iodine  vapor  into  silver  iodide,  d'his  is  evidently  an  example  of  mass 
action  [Z.  f.  anal.  Ch.  (1883)  22,  523;  also  Blau,  Monatsheft  17  (1896),  547]. 

Sunlight,  and  also  other  chemically  active  rays  (magnesium  light, 
phosphorus  light)  color  silver  chloride,  bromide,  and  iodide  at  first 
violet,  then  gray-black,  whereby  they  are  probably  converted  into  com- 
pounds of  the  form  Ag^X.  Pure  silver  iodide  is  rather  non-sensitive  to 
light,  but  exceedingly  sensitive  if  it  contains  silver  nitrate  or  substances 
which  can  take  it  up  (e.g.  tannin);  on  this  depends  the  application  of 
these  salts  in  photography. 

When  silver  plates,  previously  exposed  to  iodine,  bromine  or  chlorine  vapors,  and  con- 
sequently covered  with  a thin  layer  of  a silver  halide,  are  placed  in  a camera  obscura  in- 
visible images  are  produced.  Mercury  vapors  condense  on  the  spots  which  were  exposed 
to  the  light  and  visible  pictures  appear  (Daguerreotype,  1839).  This  was  the  first  photo- 
graj)hic  process  but  it  has  been  replaced  by  others:  (i)  Wet  collodion  process.  A 
glass  plate  is  covered  with  collodion  (a  .solution  of  pyroxylin  iii  an  ethereal  solution  of 
alcohol)  holding  in  solution  cadmium  iodide  and  ammonium  iodide  together  with  about 
one-fourth  j)art  of  ammonium  bromide.  After  the  evaporation  of  the  ether  the  glass 
plate  is  immersed  in  a solution  of  silver  nitrate,  whereby  silver  iodide  and  bromide  are 
])recipitated  uj)on  the  surface.  I'he  j)late  thus  prepared  is  expo.sed  to  light  in  the  camera 
obscura,  and,  after  the  action,  dipi)ed  into  a .solution  of  pyrogallic  acid  or  ferrous  sul- 
phate. 'fhese  reducing  substances  separate metallic  silver  in  a finely  divided  state,  which 


SILVER. 


341 


is  precipitated  upon  the  places  where  the  light  has  acted  and  the  picture,  before  invisible, 
is  “ developed.”  The  plate  is  now  introduced  into  a solution  of  potassium  cyanide  or 
sodium  hyposulphite,  which  dissolves  the  silver  salts  not  affected  by  the  light,  while  the 
metallic  unaltered  silver  remains  (fixing).  (2)  Broviide-gelati}ie />rocess.  The  plates  are 
coated  with  silver  bromide,  finely  distributed  in  the  molten  gelatine.  When  the  layers 
have  solidified  the  dry  plates  can  be  preserved  for  a long  time.  The  action  of  the  light 
produces  silver  sub-bromide  which  is  more  rapidly  reduced  by  alkaline  pyrogallol,  hydro- 
quinone,  hydroxylamine,  by  potassium  ferrous  oxalate  and  eikonogen  (amido-/?-naphthol- 
/^-monosulphonate  of  sodium)  than  unaltered  silver  bromide,  which  in  this  case  is  also 
removed  by  sodium  hyposulphite.  The  negative  thus  formed  is  covered  at  the  places 
upon  which  the  light  shone  by  a dark  layer  of  silver,  while  the  places  corresponding  to 
shadows  of  the  received  image  are  transparent.  The  copying  of  the  glass  negative  on 
paper  sensitized  with  silver  nitrate  is  executed  in  a similar  manner. 

Silver  Nitrate,  (A rgenfum  nitricmn'),  is  obtained  by  dissolving 

pure  silver  in  dilute  nitric  acid,  and  crystallizes  from  its  aqueous  solution 
in  large  rhombic  tables,  isomorphous  with  potassium  saltpeter.  At 
ordinary  temperatures  it  is  soluble  in  one-half  part  of  water  or  in  four 
parts  of  alcohol,  the  solution  has  a neutral  reaction,  and  in  this  respect 
differs  from  the  salts  of  almost  all  metals,  which  react  acid  (p.  330). 
It  fuses  at  200°,  and  solidifies  to  a crystalline  mass.  When  perfectly  pure 
it  is  not  affected  by  light.  Organic  substances  turn  it  black.  Silver 
nitrate  is  employed  for  cauterizing  wounds  {lunar  caustic). 

By  dissolving  work-silver  in  nitric  acid  a mixture  of  silver  and  copper  nitrates  is 
obtained.  To  separate  the  silver  salt  from  such  a mixture  it  is  heated  to  redness,  the 
copper  being  thus  converted  into  oxide  and  the  unaltered  silver  nitrate  extracted  with 
water. 

Silver  Nitrite,  AgN02,  is  precipitated  from  concentrated  silver  nitrate 
solutions  by  potassium  nitrite.  It  crystallizes  in  needles,  dissolves  with 
difficulty  in  water,  and  decomposes  above  90°. 

Silver  Sulphate,  Ag.^SO^,  is  obtained  by  the  solution  of  silver  in  hot  sulphuric  acid, 
and  crystallizes  in  small  rhombic  prisms  which  are  difficultly  soluble  in  water.  It  is 
isomorphous  with  anhydrous  sodium  sulphate. 

Silver  Sulphite,  AggSOg,  is  precipitated  as  a white,  curdy  mass,  if  sulphurous  acid 
be  added  to  the  solution  of  the  nitrate.  It  blackens  in  the  light  and  decomposes  at  100°: 

2 Ag2S03  = Ag^SO^  A Ag2  + SO.2. 

Silver  Nitride,  AgNg,  is  very  similar  to  silver  chloride.  It  is  more  stable  towards 
light  and  when  heated  or  struck  explodes  with  great  violence.  Angeli  claims  that  it  can 
be  readily  obtained  by  adding  a saturated  aqueous  solution  of  hydrazine  sulphate  to  a 
concentrated  aqueous  solution  of  silver  nitrite  (Ber.  26  (1893)  ill,  885  ; see  pp.  132,  324). 

Silver  Sulphide,  Ag2S,  occurs  in  regular  octahedra,  as  argentite. 
Hydrogen  sulphide  precipitates  it  as  a black  amorphous  sulphide  from 
silver  solutions.  By  careful  ignition  in  the  air  it  is  oxidized  to  silver 
sulphate.  It  is  insoluble  in  water  and  ammonium  hydroxide. 

Silver  Disulphide,  Ag2S2,  is  a dark-brown  amoridious  powder.  It  is  produced  on 
mixing  a .solution  of  silver  nitrate  in  benzonitrile  with  one  of  sulphur  in  carbon  bisulphide. 
Dilute  hydrochloric  acid  converts  it  into  a mixture  of  sulphur  and  silver  chloride,  with 
evolution  of  hydrogen  sulphide  : 

I 2IICI  = zAgCl  -f  H2S  -]-  S. 


342 


INORGANIC  CHEMISTRY. 


Silvering. — When  silver  contains  more  llian  15  per  cent,  of  copper  it  has  a yellowish 
color.  To  impart  a pure  white  color  to  objects  made  of  such  silver  they  arc  healed  to 
redness  with  acce.ss  of  air.  'I'he  copper  is  thus  superticially  oxidized,  and  may  be  removed 
by  dilute  sulphuric  acid.  'I'he  surface  of  pure  silver  is  then  polished. 

'I'he  silvering  of  metals  and  alloys  ((German  silver,  argentan ) is  executed  in  a dry  or 
wet  way.  In  the  first,  the  objects  to  be  silvered  are  coated  with  liipiid  silver  amalgam, 
with  a brush,  and  then  heated  in  an  oven  ; the  mercury  is  volatilized,  and  the  silver 
surface  then  ]X)lished. 

At  present,  the  galvanic  process  has  almost  comjdetely  superseded  the  other  i)rocesses. 
It  depends  on  the  electrolysis  of  the  .solution  of  the  double  cyanide  of  silver  and  potas- 
sium (AgCN.KCN),  whereby  the  silver  is  thrown  out  at  the  electro-negative  pole  and 
deposits  upon  the  metallic  surface  in  connection  with  that  electrode. 

'I'o  silver  glass,  cover  it  with  a mixture  of  anammoniacal  .silver  .solution,  with  reducing 
organic  substances  like  aldehyde,  milk-sugar,  and  tartaric  acid.  Under  definite  condi- 
tions, the  reduced  silver  deposits  upon  the  gla.ss  as  a regular  metallic  mirror. 


Recognition  of  Silver  Compounds. — Hydrochloric  acid  throws 
down  a white,  curdy  precipitate  of  silver  chloride,  which  dissolves  readily 
in  ammonium  hydroxide.  Zinc,  iron,  copper,  and  mercury  throw  out 
metallic  silver  from  solutions  of  silver  salts,  and  from  many  insoluble 
compounds,  like  the  chloride. 


GOLD. 

Au  = 197.2. 

Gold  {auriini)  usually  occurs  in  the  native  state,  and  is  found  dissemi- 
nated in  veins  in  some  of  the  oldest  rocks.  Gold  sands  are  formed  by  the 
breaking  and  disintegration  of  these.  It  occurs,  in  slight  quantity,  in 
the  sand  of  almost  every  river.  Combined  with  tellurium  it  forms  sylvan- 
ite,  found  in  Transylvania  and  California.  It  is  present  in  minute  quan- 
tity in  most  varieties  of  pyrites  and  in  many  lead  ores.  For  the  sepa- 
ration of  the  gold  grains  the  sand  or  pulverized  rocks  are  washed  with 
running  water,  which  removes  the  lighter  particles  and  leaves  the  specific- 
ally heavier  gold. 

The  method  of  MacArthur  and  Forrest  has  of  late  found  extensive  appli- 
cation in  the  extraction  of  gold  from  its  ores.  The  latter  after  reduction 
to  a coarse  powder  and  roasting  are  extracted  with  potassium  cyanide 
solution.  The  gold  dissolves  as  potassium  aurocyanide  and  is  precipitated 
with  metallic  zinc : 

2Au  + 4KCN  + H.p  -f  O = 2KAu(CN).,  + 2KHO. 

The  following  localities  are  important  gold  producers:  United  States, 
Australia,  Russia,  South  Africa,  and  the  Klondyke  in  Alaska. 

Native  gold  almost  invariably  contains  silver,  copper,  and  various  other 
metals.  To  remove  these,  the  gold  is  boiled  with  nitric  or  concentrated 
suliihuric  acid.  The  removal  of  the  silver  by  the  latter  acid  is  only  com- 
I)lete  if  that  metal  jiredominates ; in  the  reverse  case  a portion  of  it  will 
remain  with  the  gold.  Therefore,  to  sejiarate  pure  gold  from  alloys  poor 
in  silver  they  must  first  be  fused  with  about  three-fourths  their  weight  of 


GOLD. 


343 


the  latter  metal.  Gold  may  be  separated  from  coi)per  and  lead  by  cnpel- 
lation  (p.  337). 

Pure  gold  is  rather  soft  (almost  like  lead)  and  has  a specific  gravity  of 
19  32.  It  is  the  most  ductile  of  all  metals,  and  may  be  drawn  out  into 
extremely  fine  wire  and  beaten  into  thin  leaves,  which  transmit  green 
light.  At  about  1060°  it  meltsto  a greenish  liquid.  It  is  not  altered  by 
oxygen,  even  upon  ignition  ; acids  do  not  attack  it.  It  is  only  in  a mix- 
ture of  nitric  and  hydrochloric  acids  (aqua  regia),  which  yields  free 
chlorine,  that  it  dissolves  to  gold  chloride,  AuClg  Free  chlorine  pro- 
duces the  same.  Most  metals,  and  many  reducing  agents  (ferrous sulphate, 
oxalic  acid)  precipitate  gold  from  its  solution  as  a dark-brown  powder. 

As  gold  is  very  soft  it  wears  away  rapidly,  and  is,  therefore,  in  its  prac- 
tical applications,  usually  alloyed  with  silver  or  copper,  which  have 
greater  hardness.  The  alloys  with  copper  have  a reddish  color,  those 
with  silver  are  paler  than  pure  gold.  The  German,  French,  and  Ameri- 
can gold  coins  contain  90  per  cent,  of  gold  and  10  per  cent,  of  copper. 
A 14-karat  gold  is  generally  employed  for  ornamental  objects;  this  con- 
tains about  58.3  per  cent,  of  pure  gold  (24  karats  representing  pure  gold). 


Gold,  according  to  its  atomic  weight,  belongs  to  the  group  of  copper 
and  silver;  and,  on  the  other  hand,  forms  the  transition  from  platinum 
to  mercury.  Its  character  is  determined  to  a high  degree  by  these  double 
relations  (p.  328).  Like  the  other  elements  of  high  atomic  weight,  mer- 
cury, thallium,  lead,  and  bismuth,  belonging  to  the  same  series  of  the  pe- 
riodic system,  it  varies  considerably  in  character  from  its  lower  analogues. 

Gold,  like  silver  and  copper,  yields  compounds  of  the  form  AuX, 
aurous,  analogous  to  the  cuprous  and  argentous.  Besides,  it  has  those  of 
the  form  AuXg,  auric  derivatives,  in  which  it  is  trivalent.  These  show 
the  typical  character  of  the  trivalent  combination  form,  which  expresses 
itself  in  the  acidity  of  the  hydroxides  (p.  330);  auric  hydroxide, 
Au(OH)3,  unites  almost  solely  with  bases.  On  the  other  hand,  they 
show  many  similarities  to  the  highest  combination  forms  of  the  metals 
with  high  atomic  weight:  platinum  (PtXJ,  mercury  (HgX2),  thallium 
(TIX3),  and  lead  (PbXJ  (p.  357). 


AUROUS  COMPOUNDS. 

Aurous  Chloride,  AuCl,  is  produced  by  heating  auric  chloride, 
AuClg,  to  180°,  and  forms  a white  powder  insoluble  in  water.  When 
ignited,  it  decomposes  into  gold  and  chlorine;  cold  water  decomposes  it 
slowly,  more  quickly  on  heating  with  formation  of  auric  chloride  and  gold. 

Aurous  Iodide,  Aul,  separates  as  a yellow  powder,  if  potassium 
iodide  be  added  to  a solution  of  auric  chloride  : 

AUCI3  + 3KI  = Aul  T L 4-  3KCI. 

When  heated  it  breaks  up  into  gold  and  iodine. 


344 


INORGANIC  CHEMISTRY. 


When  auric  oxide  or  sulpliicle  is  dissolved  in  potassium  cyanicU',  largo 
colorless  prisms  of  the  double  cyanide,  AuCN.KCN,  crystallize  out  iijx)!! 
evaporation.  The  galvanic  current  and  many  metals  ])recipitate  gold 
from  this  compound  ; hence  it  serves  for  the  sej)aration  of  gold  from 
ores  and  for  electrolytic  gilding,  which,  at  })rescnt,  has  almost  entirely 
superseded  the  gilding  in  the  dry  way  (see  ]).  342). 

Aurous  Oxide,  Au^O,  is  formed  by  the  acticm  of  potassium  hydroxide 
upon  aurous  chloride.  It  is  a dark-violet  j)owder  which  at  250°  decom- 
poses into  gold  and  oxygen.  It  is  changed  to  auric  chloride  and  gold 
by  the  action  of  hydrochloric  acid. 

Only  a few  double  salts  of  the  oxygen  derivatives  of  univalent  gold  are 
known. 


AURIC  COMPOUNDS. 

Auric  Chloride,  AuClg,  results  from  the  action  of  chlorine  upon  the 
metal.  It  is  a lemon-yellow  deliquescent  mass,  which  dissolves  readily  in 
alcohol  and  in  ether.  On  evaporating  the  solution  long,  yellow-colored 
needles  of  the  composition  HAuCl^.  qli/),  hydrochlor-auric  acid,  remain. 
It  forms  beautifully  crystallized  double  salts  with  many  metallic  chlorides, 
g.,  KAuCl^  -}-  2H2O  and  NH^AuCl^  -f-  2^H20.  When  a solution  of 
auric  chloride  is  heated  with  magnesium  oxide  a brown  precipitate  is 
obtained,  from  which  all  the  magnesia  is  removed  by  concentrated  nitric 
acid,  leaving  Auric  Oxide  (AiqOg).  This  is  a brown  powder  which 
decomposes,  near  250°,  into  gold  and  oxygen.  If  the  preci})itate  con- 
taining the  magnesia  be  treated,  not  with  concentrated,  but  with  dilute 
nitric  acid,  Auric  Hydroxide,  Au(OH)g,  remains  as  a yellowish-red 
powder:  Both  the  oxide  and  hydroxide  are  insoluble  in  water  and  acids; 
they  possess,  however,  acid  properties,  and  dissolve  in  alkalies.  There- 
fore the  hydroxide  is  also  called  auric  acid.  Its  salts,  the  aurates,  are 

I 

constituted  according  to  the  formula  MeAuOg,  and  are  derived  from  the 
meta-acid,  HAuOg  = AuO.  OH. 

Potassium  Aurate,  KAuOg  + 3H2O,  crystallizes  in  bright  yellow 
needles,  from  a potassium  hydroxide  solution  of  auric  oxide.  These  are 
readily  soluble  in  water  ; the  solution  reacts  alkaline.  The  corresponding 
aurates  are  precipitated  from  this  solution  by  many  metallic  salts,  e.  g.  ; 


KAuOg  + AgNOg  = AgAuO^  + KNOg. 

The  precipitate  produced  by  magnesia  in  a solution  of  auric  chloride 
(see  above)  consists  of  magnesium  aurate,  Mg(Au02)2.  ’ Oxygen  salts  of 
auric  oxide  are  not  known. 

Auric  Sulphide,  AiqSg,  is  preci])itated  as  a blackish-brown  com- 
]>ound,  from  gold  solutions,  by  liydrogen  sulphide.  Its  composition  is 
either  AiqSg  or  AU2S2  (mixed  with  sulphur),  depending  on  the  condi- 
tions prevailing  at  the  time.  It  dissolves  in  alkaline  sulphides  with  forma- 
tion of  sulj)ho-salts,  which  are  only  derived  from  Au2S(<f.  ^.,  KgAuSj). 
Hydrogen  sulphide  })reci})itates  aurous  sulphide,  AiqS,  from  hot  gold 


METALS  OF  GROUP  III.  345 

solutions.  It  is  steel-gray  in  color,  and  soluble  in  pure  water,  from 
which  hydrochloric  acid  precipitates  it.  It  is  used  in  gold  plating. 

Stannous  chloride  (SnClJ  added  to  an  auric  chloride  solution  pro- 
duces, under  certain  conditions,  a purple-brown  precipitate,  purple  of 
Cassius,  which  is  employed  in  glass  and  porcelain  painting.  Alumina 
and  magnesia  yield  similar  purples,  and  it  appears  that  their  red  colora- 
tion is  due  to  finely  divided  metallic  gold. 

On  pouring  ammonium  hydroxide  over  auric  oxide  a brown  compound 
is  produced — fubninating  gold.  When  this  is  dried  and  heated  or  struck 
it  explodes  very  violently. 


METALS  OF  GROUP  III. 

The  trivalent  elements,  affording  derivatives  mainly  of  the  form  MeX3, 
belong  to  group  111  of  the  periodic  system  (p.  246)  : 

Sc  = 44.1  Y 89  La  = 138  Yb  = 173 

B =:r  II  A1  = 27.1 

Ga  =go  In  ^-114  ...  T1  = 203.7. 

These  bear  the  same  relation  to  one  another  as  do  the  elements  of 
group  II  (p.  299).  Boron  has  the  lowest  atomic  weight,  and  the  basic, 
metallic  character  in  it  is  reduced  very  much  or  does  not  appear  at  all. 
In  its  exclusively  acidic  hydroxide,  B(OH)3,  it  approaches  the  metal- 
loids, and  is  therefore  treated  with  them  (p.  241). 

Aluminium  is  a perfect  metal;  its  hydroxide,  Al(OH)3,  exhibits  a pre- 
dominating basic  character,  and  yields  salts  with  acids.  Its  relations  to 
boron  are  like  those  of  silicon  to  carbon,  or  of  magnesium  to  beryllium. 
The  connection  of  aluminium  and  boron  with  the  same  group  plainly 
shows  itself  in  the  entire  character  of  the  free  elements,  and  in  their 
compounds.  Thus  aluminium  and  boron  are  not  dissolved  by  nitric  acid, 
but  by  boiling  alkalies  : 

A1  -p  3KOH  = A1(0K)3  + 3H. 

There  is  only  a gradual  difference  between  their  hydrates.  Boron 
hydroxide,  B(OH)3,  not  only  acts  as  a feeble  acid,  but  we  also  find  that 
aluminium  hydroxide  manifests  an  acidic  character,  inasmuch  as  it  is 
capable  (pp.  258,  330)  of  forming  metallic  salts  with  strong  bases  (chiefly 
the  alkalies) ; but  owing  to  the  higher  atomic  weight  of  aluminium  the 
basic  character  exceeds  the  acidic. 

Scandium,  yttrium,  lanthanum  and  ytterbium  attach  themselves  to 
aluminium  as  the  first  sub-group.  These  constitute  the  third  members  of 
the  great  periods,  and  hence  exhibit  a pronounced  basic  character.  As 
light  metals,  they  are  very  similar  to  aluminium  in  their  compounds,  so 
that  they  are  all  embraced  in  one  group,  which  (corresponding  to  the 
earthy  nature  of  their  oxides)  is  designated  the  group  of  earth  metals. 
Cerium,  praseodymium  and  neodymium  bear  a peculiar  relation  to  lan- 
thanum ; their  atomic  weights  are  nearly  alike  and  their  properties  very 


346 


INORGANIC  CHEMISTRY. 


similar.  They,  like  erbium,  terbium,  llmlium,  holmium  aud  samarium, 
are  by  no  means  so  accurately  known  that  they  can  l)e  assigned  definite 
])laces  in  the  periodic  system.  Most  of  them  are  i)robably  mixtures  of 
unknown  elements. 

The  second  sub  group  is  more  distinctly  characterized  and  accurately 
investigated;  it  consists  of  the  heavy  metals,  gallium,  indium  and  thal- 
lium. 'I'hese  belong  to  the  right  side  of  the  great  periods,  ])ossess,  there- 
fore, a less  basic  character,  and  bear  the  same  relation  to  one  another  as 
zinc,  cadmium  and  mercury. 


I.  GROUP  OF  THE  EARTH  METALS. 

ALUMINIUM. 

A1  = 27.1. 

Aluminium,  both  in  quantity  and  distribution,  is  one  of  the  most 
important  constituents  of  the  earth’s  crust.  As  oxide,  it  crystallizes  as 
ruby,  sapphire  and  corundum ; less  pure  as  emery.  It  is  commonly 
found  as  aluminium  silicate  (clay,  kaolin),  and  in  combination  with  other 
silicates,  as  felds])ar,  mica,  and  also  in  most  crystalline  rocks.  It  occurs, 
too,  united  with  fluorine  and  sodium,  as  cryolite,  AlFl3.3NaFl,  in  large 
deposits,  in  Greenland. 

Metallic  aluminium  is  at  present  obtained  almost  exclusively  by  the 
electrolysis  of  alumina  dissolved  in  fused  cryolite.  The  oxygen  burns 
the  carbon  anode  to  carbon  monoxide  and  aluminium  separates  in  the 
molten  condition. 

Metallic  aluminium  was  first  obtained  as  a gray  powder  by  Wohler  in  1827  when  he 
heated  the  chloride  with  potassium.  Later  he  conducted  the  vapors  of  aluminium  chlo- 
ride over  heated  sodium  or  potassium,  obtaining  thereby  globules  of  the  metal.  Bunsen, 
1854,  reduced  the  double  chloride  of  sodium  and  aluminium  with  the  electric  current  and 
separated  aluminium  also  in  powder  form.  St.  Claire-Deville  announced  similar  methods 
almost  simultaneously.  It  is  due  to  his  efforts  that  the  metal  was  produced  on  a com- 
mercial scale  (1856).  This  he  accomplished  by  reducing  aluminium  sodium  chloride, 
partly  with  introduction  of  cryolite,  by  means  of  sodium  : 

AlCb.NaCl  + 3Na  = 4NaCl  -f  Al. 

However,  in  1888  this  method  was  in  use  in  but  one  factory.  A new  impulse  was  given 
to  the  manufacture  of  the  metal  after  the  di.scovery  of  Ca.stner’s  and  of  Netto’s  methods  of 
preparing  metallic  .sodium.  Equal  parts  of  cryolite  and  sodium  chloride  were  fused 
together  and  blocks  of  sodium  of  5-7  kilograms  in  weight  were  dipped  into  the  then 
molten  mass,  the  slag,  separated  from  the  aluminium  and  containing  .sodium  fluoride,  was 
treated  with  aluminium  sul[ihate  and  artificial  cryolite  prepared.  This  sodium  method 
was  abandoned  as  a conse(|uence  of  the  develojmient  of  electrolysis.  In  America 
natural  gas  is  used  for  heat  and  jiower  purpo.ses  ; and  the  power  of  Niagara  Falls  was 
turned  to  the  production  of  electric  currents.  A portion  of  the  Rhine  Falls  at  Schaff- 
liauscn  in  Switzerland  is  applied  to  the  same  purpo.se.  Another  large  aluminium  plant 
exists  at  hroges  near  Orenoble.  'I'he  electrolysis  is  conducted  in  accordance  with  the 
method  of  1 Itroult,  using  a .solution  of  alumina  in  fused  cryolite  and  fluorspar.  T he  oxide 
breaks  down  (juite  readily  into  aluminium  and  oxygen.  The  Hall  method,  so  largely 


Al^UiMINlUM. 


347 


used  in  America,  is  practically  the  same.  (Compare  Cl.  Winkler,  Ch.  Zt.,  1892,  349, 
and  Elbs;  ibid.,  1894,  1637;  1897,  995;  especially  W.  Borcher’s  Elektrometallurgie, 
1896,  and  his  Elektrische  Schmelzofen,  Halle,  1897.) 

Aluminium  is  an  almost  silver-white  metal  of  strong  luster,  very  ductile, 
and  it  may  be  drawn  out  into  fine  wire  and  beaten  into  thin  leaflets.  Its 
specific  gravity  is  2.583  ; it  belongs,  consequently,  to  the  light  metals  (see 
pp.  252,313).  It  fuses  at  660°  but  will  not  vaporize.  It  changes  very  little 
in  the  air  at  ordinary  temperatures,  and  even  when  heated.  If,  however, 
thin  leaves  be  heated  in  a stream  of  oxygen,  they  will  burn  with  a bright 
light.  Pure  aluminium  is  attacked  with  difficulty  by  water.  The  commer- 
cial article,  containing  sodium,  is  more  readily  affected,  while  the  amal- 
gamated product  is  very  easily  and  quickly  acted  upon  by  the  reagent. 
Nitric  acid  affects  aluminium  only  superficially  ; sulphuric  acid  dissolves 
it  on  boiling,  while  it  is  readily  soluble,  even  in  the  cold,  in  hydrochloric 
acid,  and  in  a solution  containing  5 per  cent,  of  acetic  acid  and  some 
sodium  chloride  it  is  attacked  at  once.  It  dissolves  in  potassium  and 
sodium  hydroxide  with  evolution  of  hydrogen,  and  forms  aluminates: 

A1  4-  3KOH  = K3  AIO3  -f  3H. 

Owing  to  its  stability  in  the  air  and  beautiful  luster,  aluminium  is  some- 
times employed  for  vessels  and  ornaments. 

It  has  been  found  that  vessels  of  aluminium  are  attacked  by  most  foods  and  drinks. 
The  effect  is  slight  and  diminishes  rapidly  with  continual  use.  The  relish  of  foods  pre- 
pared in  such  vessels  appears  to  suffer  nothing  by  continuance  nor  does  there  appear  to 
be  any  harmful  influences  to  the  general  health  (Plagge). 

Alloys  of  copper  with  5-12  per  cent,  of  aluminium  are  distinguished 
by  their  great  hardness  and  durability.  They  make  good  castings,  and 
possess  a gold-like  color  and  luster.  Under  the  name  of  aluminium 
bronze,  they  are  used  for  the  composition  of  various  articles,  as  watches, 
spoons,  etc.  Their  firmness  and  elasticity  render  them  suitable  for  physical 
instruments  (arms  of  balances)  and  watch  springs. 


Aluminium  has  a most  remarkable  affinity  for  oxygen.  It  has  conse- 
quently been  in  use  for  quite  a while  in  isolating  other  elements  from  their 
oxides  (p.  136).  Magnesium  oxide  is  the  only  oxide  not  reduced  by 
metallic  aluminium.  Alumina  (corundum)  always  results  in  such  reduc- 
tions : 

Cr203  -p  2AI  = AI2O3  + Cr2. 

When  metallic  oxides  are  heated  with  aluminium  powder  so  much  heat  is  liberated 
that  violent  explosions  take  place.  II.  Gold.schmidt  has  recently  discovered  a means  of 
controlling  these  rearrangements  and  of  utilizing  the  immense  heat  of  the  reaction  by  ignit- 
ing the  mixture  of  metallic  oxide  and  aluminium  powder  within  by  means  of  a percussion 
point  at  but  one  place  and  letting  the  transpositions  advance  from  there.  The  “ percus- 
sion point”  consists  of  a magnesium  band  supplied  at  its  end  with  a mixture  of  aluminium 
powder  and  barium  peroxide  or  some  other  sub.stance  rich  in  oxygen.  In  such  reactions 
the  aluminium  is  converted  into  a slag  of  fused  corundum,  which  on  cooling  crystallizes 
in  part  like  the  natural  product. 


348 


INORGANIC  CHKMISTRV. 


Sulphides  iiisiciul  of  oxides,  can  be  used  ; tliis  is  sonietitnes  advantageous  because 
aluuiiniuin  sulphide  melts  easily  and  can  be  readily  separated  from  melals  : 

3CuS  4 2AI  = AlgSg  -f  3CU. 

[See  II.  Goldschmidt,  Ann.  Chem.  (1898)  301,  19.] 

d'he  heat  evolved  in  the  formation  of  aluminium  hydroxide  is  388.8  Cal.  This  far 
exceeds  that  of  any  other  oxides  : 

Cu,0  = 37  ; Fe„03,3llg0  = 191. 

ddie  heat  corresponding  to  one  atom  of  oxygen,  is,  therefore,  129.6  Cal.;  since  that  of 
water  is  far  less  (11.2,0  — 69.0),  it  must  be  decompo.sed  by  aluminium,  with  liberation 
of  hydrogen  (p.  273).  If  this  does  not  transpire  under  ordinary  conditions,  the  reason 
must  be  sought  for  in  the  insolubility  of  aluminium  hydroxide.  Indeed,  the  reaction 
occurs  if  aluminium  chloride,  or  another  salt,  in  which  the  aluminium  oxide  is  soluble, 
be  added  to  the  water.  Conversely,  the  high  heat  of  formation  of  aluminium  oxide 
explains  why  it  is  not  reduced  by  carbon. 


Aluminium  unites  in  but  one  proportion  with  oxygen,  Al^Og;  this  is 
the  source  of  salts  of  the  type  AIX^  (p.  349).  Water  decomposes  them 
hydrolytically  in  part  and  that  explains  the  acid  reaction  of  their  aqueous 
solutions;  they  have  a sweet,  acrid  taste. 

The  trivalence  of  aluminium  is  shown  in  its  organic  derivatives,  e.g.  : 

A1(CH3)3,  A\{C,H,)„  etc. 

Such  compounds  like  those  with  hydrogen  are  best  adapted  for  the  determination  of  the 
valence  of  an  element  (p.  247).  At  sufficiently  high  temperatures  the  density  of  the 
vapors  of- aluminium  chloride  correspond  to  the  formula  AlCl,,  but  at  lower  temperatures 
this,  like  other  metallic  chlorides  (<?.  g.,  tin  dichloride,  gallium  chloride,  etc.,  p.  339), 
shows  a tendency  to  form  larger  molecules — to  polymerize. 


Aluminium  Chloride,  AICI3,  is  produced  by  the  action  of  chlorine 
upon  heated  aluminium;  also  by  heating  a mixture  of  aluminium  oxide 
and  carbon  in  a current  of  chlorine: 

AlA  + 3C  -p  6C1  = 2AICI3  + 3CO. 

Chlorine  and  carbon  do  not  act  separately  upon  the  oxide  ; by  their 
mutual  action,  however,  the  reaction  occurs  in  consequence  of  the  affinity 
of  carbon  for  oxygen,  and  of  chlorine  for  aluminium.  The  oxides  of 
boron  and  silicon  sliow  a similar  deportment. 

In  order  to  get  chloride  free  from  iron  and  oxychloride,  it  is  advisable 
to  heat  the  metal  in  dry  hydrochloric  acid  gas  (Seubert  and  Pollard,  Ber. 
(1891)  ?4,  2575). 

Aluminium  chloride  may  be  obtained  in  white,  hexagonal  leaflets  by 
sublimation.  It  sublimes  readily,  but  will  only  fuse  when  subjected  to 
high  ])ressure  (in  a sealed  tube)  and  at  193-194°.  At  175-179°  the 
fused,  cooling  mass  exhibits  the  jihenomenon  of  boiling  and  subliming 
(j).  229).  The  boiling  point  under  the  ordinary  pressure  is  183°. 


ALUMINIUM. 


349 


Deville  and  Troost  (1857)  first  determined  the  vapor  density  of  aluminium 
chloride,  bromide,  and  iodide  at  440°,  and  found  it  to  correspond  to  the 
formulas  Al2Clg,  Al2Brg,  and  Later,  Friedel  and  Crafts  redetermined 

that  of  the  chloride  at  218-440°  with  like  results.  This  led  to  the  sup- 
position that  aluminium  was  quadrivalent  in  all  its  compounds  (Jahresb. 
f.  Ch.  1888,  I,  131).  The  most  recent  investigations  of  Nilson  and 
Peterson  (Z.  f.  phys.  Ch.  i (1887),  459)  have  shown  that  at  higher 
temperatures  (above  760°)  the  vapor  density  of  the  chloride  corresponds 
to  the  formula  AICI3  (p.  339). 

Aluminium  chloride  absorbs  moisture  from  the  air,  and  deliquesces. 
It  crystallizes  from  a concentrated  hydrochloric  acid  solution,  with  six 
molecules  of  water.  On  evaporating  the  aqueous  solution,  the  chloride 
decomposes  into  aluminium  oxide  or  hydroxide  and  hydrogen  chloride : 

2AICI3  + 3H2O  = AI2O3  + 6HC1. 

It  forms  double  chlorides  with  many  metallic  chlorides,  viz.  : AICI3.  NaCl, 
AICI3.KCI.  The  solutions  of  these  may  be  evaporated  to  dryness  with- 
out decomposition.  It  also  unites  with  many  halogen  derivatives  of  the 
metalloids : 

A1C13 . PCI.,  AICI3 . POClg,  AICI3 . SCh. 

Aluminium  Bromide,  AlBrj,  is  obtained  like  the  chloride,  and  consists  of  shining 
leaflets  which  fuse  at  90°  and  boil  at  265-270°. 

Aluminium  Iodide,  AII3,  is  formed  on  heating  aluminium  filings  with  iodine.  It 
is  a white,  crystalline  mass,  fusing  at  185°,  and  boiling  at  360°.  It  is  best  prepared  by 
covering  sheet  aluminium  with  carbon  bisulphide,  and  then  adding  the  calculated  amount 
of  iodine  gradually,  letting  the  whole  stand  for  some  time,  and  then  distilling  off  the  car- 
bon bisulphide.  The  reaction  occurring  between  aluminium  iodide  and  oxygen  is  inter- 
esting. If  the  vapor  of  the  former  be  mixed  with  the  latter,  and  then  brought  in  contact 
with  a flame,  or  if  acted  upon  by  an  electric  spark,  a violent  detonation  will  ensue  ; 
aluminium  oxide  and  iodine  result : 

2AII3  3O  = AI2O3  -f-  312- 

This  deportment  is  due  to  the  great  difference  in  the  heats  of  formation  of  the  aluminium 
oxide  (about  380  Cal.),  and  the  iodide  (140.6  Cal.).  The  chloride  and  bromide  are 
similarly  decomposed,  but  with  less  violence. 

Aluminium  Fluoride,  AIFI3,  obtained  by  conducting  hydrogen  fluoride  over  heated 
aluminium  oxide  or  hydroxide,  sublimes  at  a red  heat  in  colorless  rhombohedra.  It  is 
insoluble  in  water,  unaltered  by  acids,  and  is  very  stable.  It  yields  insoluble  double 
fluorides  with  alkaline  fluorides.  The  compound,  AlFl3.3NaFl,  occurs  in  Greenland, 
in  large  deposits,  as  cryolite,  and  is  employed  in  the  soda  and  alumina  manufacture 
(p.  292). 

Aluminium  Oxide,  AI2O3,  is  found  crystallized  in  hexagonal  prisms 
in  nature,  as  ruby,  sapphire,  and  corundum,  colored  by  other  admixtures. 
Impure  corundum,  containing  aluminium  and  iron  oxides,  is  called 
emery,  and  .serves  for  polishing  glass.  The  specific  gravity  of  these 
minerals  is  3.9  ; their  hardness  is  only  a little  below  that  of  the  diamond. 
Artificial  aluminium  oxide — corundum — may  be  obtained  by  reducing 
metallic  oxides  with  aluminium — following  Goldschmidt’s  method;  in 
making  chromium  the  slag  will  contain  little  crystals  of  the  ruby.  The 
crystallized  or  strongly  ignited  aluminium  oxide  is  almost  insoluble  in 


350 


INORGANIC  CHEMISTRY. 


acids;  to  decompose  it,  it  is  fused  with  caustic  alkalies  or  with  ])rimary 
potassium  sulphate,  flKSO^. 

Aluminium  Hydroxides. — The  fior?nal hydroxide,  Al(OH)3,  occurs 
in  nature  as  hydrargillite.  The  hydroxide,  Al202(()H)^,  is  diaspore. 
Bauxite  is  a mixture  of  the  hydroxide,  Al20((i)H)^,  with  ferric  oxide.  It 
occurs  in  large  dei)osits  in  France,  in  the  Caucasus,  etc.  The  normal 
hydroxide  is  artificially  obtained  as  a white  voluminous  i)recipitate, 
by  adding  ammonium  hydroxide  or  an  alkaline  carbonate  (in  the  latter 
case  carbon  dioxide  escapes,  p.  352)  to  a soluble  aluminium  salt.  It 
dissolves,  when  freshly  precipitated,  in  acids  and  in  potassium  and 
sodium  hydroxides,  liy  long  standing  under  water,  or  after  drying, 
it  is,  without  any  alteration  in  comi)osition,  difficultly  soluble  in 
acids.  When  carefully  heated,  the  normal  hydroxide  first  passes  into 
the  hydrate,  AlO.OH. 

The  freshly  preeijiitated  hydroxide  dissolves  readily  in  a solution  of  aluminium  chloride 
or  acetate.  On  dialyzing  (p.  237)  this  solution  the  aluminium  salt  or  crystalloid  diffuses, 
and  in  the  dialyzer  remains  the  pure  acjueous  solution  of  the  hydroxide.  This  has  a faint 
alkaline  reaction  and  is  coagulated  by  slight  quantities  of  acids,  alkalies  and  many  salts  ; 
the  soluble  hydrate  passes  into  the  insoluble  gelatinous  modification. 

Gelatinous  aluminium  hydroxide  possesses  the  property  of  precipitating  many  dyestuffs 
from  their  solutions,  forming  colored  insoluble  compounds  (lakes)  with  them.  On  this 
is  based  the  application  of  aluminium  hydroxide  as  a mordant  in  dyeing.  The  acetate  is 
generally  used  for  this  purpose.  Goods  saturated  with  this  salt  are  heated  with  steam, 
which  causes  the  decomposition  of  the  weak  acetate  ; acetic  acid  is  driven  off,  while 
the  separated  aluminium  hydroxide  adheres  to  the  fiber  of  the  material.  If  the  latter 
now  be  introduced  into  the  solution  of  coloring  matter  the  color  will  be  fixed  by  the 
aluminium  hydroxide  upon  the  fiber.  At  present,  sodium  aluminate  is  employed  in- 
stead of  the  acetate. 

Aluminium  hydroxide  has  a feeble  acid  character,  and  can  form  salt- 
like compounds  with  strong  bases.  On  carefully  evaporating  its  solution 
in  sodium  or  potassium  hydroxide,  or  upon  addition  of  alcohol,  white 
amorphous  compounds  of  KAIO2,  NaA102  and  Al(ONa)3  are  obtained. 
The  potassium  compound  can  be  obtained  in  crystalline  form.  These 
derivatives,  known  as  aliwiinates,  are  not  very  stable,  and  are  even 
decomposed  by  carbon  dioxide,  with  elimination  of  aluminium  hydroxide : 

2NaA10.,  4-  CO2  + 3H,0  = 2A1(0H)3  + 

d'he  aluminium  hydroxide  obtained  in  this  manner,  in  distinction  from 
that  precipitated  from  acid  aluminium  solutions  by  the  alkalies,  is  not 
gelatinous,  and  is  more  difficultly  soluble  in  acids,  especially  acetic.  It 
comprises  the  ordinary  alumina  of  commerce. 

On  adding  calcium  chloride,  strontium  chloride,  or  barium  chloride  to 
the  solution  of  potassium  or  sodium  aluminate,  white  insoluble  aluminates 
are  precipitated  : 

2NaA102  4-  CaCl2  = Ca(A102).,  4-  2NaCl. 

Similar  aluminates  fre(piently  occur  as  crystallized  minerals,  in  nature. 
Thus  the  spinels  consist  chiefly  of  magnesium  aluminate. 


ALUMINIUM. 


351 


chrysoberyl  is  beryllium  aluminate,  gahnite  is  zinc  alu- 

. , AlO.O.  .7 

minate,  ^10.0^^^^' 

Nearly  all  these  minerals,  commonly  called  spinels,  crystallize  in  regu- 
lar octahedra,  like  the  corresponding  chromium  compounds  (see  these); 
the  exceptions  are  chrysoberyl,  crystallizing  in  the  rhombic  system,  and 

hausmannite,  = MnO'  quadratic  system. 

Technically,  alumina  is  obtained  from  cryolite,  bauxite  and  other  minerals  containing 
aluminium.  The  pulverized  bauxite  is  heated  ’with  dry  sodium  carbonate  in  furnaces, 
and  the  resulting  sodium  aluminate  extracted  with  water.  P'rom  the  clear  solution  car- 
bon dioxide  precipitates  the  hydroxide,  while  sodium  carbonate  remains  dissolved,  and  is 
afterward  recovered.  The  dried  aluminium  hydroxide  occurs  as  a white  powder  in  trade. 

K.  j.  Baeyer  treats  the  finely  ground  bauxite  with  caustic  soda  and  stirs  into  this  solu- 
tion freshly  precipitated,  crystalline  alumina  (as  obtained  on  conducting  carbon  dioxide 
into  a cold  aluminate  solution).  All  of  the  alumina  separates  after  some  hours,  and  the 
alkaline  liquid  is  again  run  through  the  same  course.  The  method  is  based  on  the  decom- 
position of  the  aluminates  by  water  ; the  alumina  which  is  mixed  in  with  it  acts  like  a 
crystal  of  Glauber  salt  upon  supersaturated  solutions  (p.  288). 

The  gelatinous,  readily  soluble  (colloidal)  aluminium  hydroxide  (see  above)  precipi- 
tated from  acid  solutions  by  alkalies,  has  lately  been  prepared  upon  a large  scale,  accord- 
ing to  the  method  of  Lowig,  by  treating  the  sodium  aluminate  solution  with  milk  of  lime  ; 
calcium  aluminate  precipitates,  while  sodium  hydroxide  remains  in  solution  : 

2NaA102  -b  Ca(0H)2  = Ca(A102)2  + 2NaOH. 

The  calcium  aluminate  is  dissolved  in  hydrochloric  acid  : 

Ca(A102)2  + 8IIC1  = 2AICI3  -f  CaCq  + 4H2O, 

and  to  the  solution  now  containing  the  alumina  as  chloride  the  corresponding  amount  of 
calcium  aluminate  added,  and  aluminium  hydroxide  is  precipitated  : 

2AICI3  4-  3Ca(A102)2  4-  12H2O  ==  8A1(0H)3  + 3CaCl2. 

According  to  this  procedure,  the  sodium  hydroxide  formed  in  the  first  reaction  is  obtained 
together  with  the  alumina. 

On  conducting  carbon  dioxide  into  a solution  of  alkaline  carbonates,  and  adding  a 
solution  of  an  alkaline  aluminate  at  the  same  time,  white  aluminiian-alkali  carbonates 
are  precipitated  : 

AI2O3.K2O  4-  2KHCO3  = Al203.K,0.2C02  4 2K0fI. 

The  caustic  alkali  that  is  formed  in  this  way  is  converted  again  into  bicarbonate  by  car- 
bon dioxide.  In  a dry  state  the  precipitates  are  white,  chalk-line  masses  which  at  90° 
contain  five  molecules  of  water  : 

Al2O3.K2O.2CO2  f 5H.A 

Their  constitution  may  be  expressed  by  the  formula : 

K-O-CO-O-AI  ^ O. 

They  dissolve  readily  in  dilute  acids,  even  acetic,  with  evolution  of  carbon  dioxide,  and 
are  suitable  for  the  preparation  of  pure  alumina  mordants  and  antiseptic  solutions  (Lowig). 


352 


INORGANIC  CHEMISTRY. 


Tlie  basic  cliaracter  of  altiniinium  liydroxide  exceeds  the  acid  ; but  it 
is  so  feeble  that  it  is  not  ca])able  of  forming  salts  with  weak  acids,  as  car- 
bon dioxide,  suliihnrons  acid,  and  hydrogen  sulphide.  When  sodium 
carbonate  is  added  to  solutions  of  aluminium  salts,  aluminium  hydroxide 
is  precipitated,  while  carbon  dioxide  is  set  free  : 

2AICI3  + sNa^COj  + 311,0  = 2AI  (011)3  I 6NaCl  + 3CO,. 


The  alkaline  sulphides  behave  similarly  : 

2AICI3  -1  3(Nnd,S  + 611,0  = 2A1(0II)3  + 6NIbCl  -|-  3H,S. 

That  the  hydrate  is  a feeble  acid  is  shown  in  the  decomposition  of  the 
alkali  salts  by  carbon  dioxide. 

Aluminium  Sulphate,  Al2(SOj3,  crystallizes  from  aqueous  solution 
with  eighteen  molecules  of  water  in  thin  leaflets  with  ])early  luster.  These 
dissolve  readily  in  water;  when  heated,  they  melt  and  lose  all  their  water 
of  crystallization.  The  sulphate  is  obtained  by  dissolving  the  hydroxide 
in  sulphuric  acid,  or  by  the  decomposition  of  pure  aluminium  silicate 
(clay)  with  the  same  acid  ; the  residual  silicic  acid  is  removed  by  filtra- 
tion, and  the  solution  of  the  sul])hate  evaporated.  When  a quantity  of 
ammonium  hydroxide,  insufficient  for  complete  precipitation,  is  added 
to  the  sulphate,  basic  sulphates  separate  out.  Salts  similar  to  the  latter 
are  also  found  in  nature;  thus,  aluminite,  used  to  prepare  alum,  has  the 
composition  : 

+ 7H3O  or  (A10.0),S0,  + 9H,0. 

Aluminium  sulphate  can  combine  with  the  alkaline  sulphates  and 
affords  double  salts,  termed  alu??is,  e.  g.,  potassium  alum: 

Al2(SOj3.K,SO,  q 24H,0  or  KAl(SOd,  + I2H,0. 

Their  constitution  is  expressed  by  the  following  formula : 

III 

SO^  = A1  — SO,  — K + i2H,0. 

In  this  compound  the  potassium  may  be  replaced  by  sodium,  ammonium, 
rubidium,  Ccesium,  and  also  by  thallium  hydroxylamine,  and  some  organic 
bases.  Iron,  chromium  and  also  vanadium,  titanium  and  manganese 
afford  like  derivatives: 

Fe,fS0,)3.K,S0,  + 24H,0  Cr,(SO,)3.  (NT-Tp,(SO),  + 2411,0. 

Potassium  iron  alum.  Ammonium  chromium  alum. 


All  these  alums  crystallize  in  regular  octahedra  or  cubes,  and  can  form 
isomorphous  mixtures. 

'I'he  most  im|)ortant  of  them  is  pofassiinn  aluminium  sulphate  or  ordi- 
nary alum,  KA1(S0,)2 -|-  12!  1,0.  Jt  crystallizes  from  water  in  large, 
transjiarent  octahedra,  soluble  in  8 parts  of  water  of  ordinary  tempera- 


ALUMINIUM. 


353 


tiire,  or  in  3<3  part  of  boiling  water.  The  solution  has  an  acid  reaction  and 
a sweetish,  astringent  taste.  When  })laced  over  sulphuric  acid,  alum  loses 
three-fourths  of  its  water.  When  heated  it  melts  in  its  water  of  crystalliza- 
tion, loses  all  the  latter,  and  becomes  a white,  voluminous  mass — burnt 
alum.  Upon  adding  a little  sodium  or  potassium  carbonate  to  a hot 
alum  solution,  the  hydroxide  first  produced  dissolves,  and  when  the  liquid 
cools,  the  alum  crystallizes  out  in  cubes,  as  cubical  alimi.  The  addition 
of  more  sodium  carbonate  causes  the  precipitation  of  the  basic  salt — 
KA1(S0^)2.  A1(0H)3.  Alunite,  found  in  large  quantities  near  Rome  and 
in  Hungary,  has  a similar  composition,  K(A\0').J SO^)^  3H2O. 


Commercial  alum  is  obtained  according  to  various  methods:  (i)  From  alunite,  by 
heating  and  extracting  with  hot  water.  In  this  way  alum  dissolves  while  the  hydroxide 
remains  ; from  such  solutions  the  former  crystallizes  in  combinations  of  the  octahedron 
with  cube  faces — Roman  alum.  (2)  The  most  common  source  of  alum  was  formerly  altan 
shale,  a clay  containing  pyrite  and  peat.  This  is  roasted  and  after  moistening  with  water 
is  exposed  for  a long  time  to  the  action  of  the  air.  By  this  means  pyrite,  FeS2,  is  con- 
verted into  ferrous  sulphate,  FeS04,  and  free  sulphuric  acid,  which,  acting  upon  the 
clay,  forms  aluminium  sulphate.  The  mass  is  extracted  with  water,  potassium  sulphate 
added,  and  the  whole  permitted  to  crystallize.  (3)  At  present  clay  is  treated  directly  with 
sulphuric  acid,  and  to  the  solution  of  aluminium  sulphate  potassium  or  ammonium  sul- 
phate is  added.  (4)  Bauxite  and  cryolite  are  admirable  material  for  the  preparation  of 
alum.  The  working  of  cryolite  for  alumina  and  soda  is  described  on  p.  292,  and  that  of 
bauxite,  p.  351). 


Ammonium  Alum,  NH4A1(S04)2  4-  12H2O,  crystallizes,  like  potas- 
sium alum,  in  large  crystals,  and  at  present,  owing  to  its  cheapness,  is 
applied  almost  exclusively  for  technical  purposes.  Sodium  alum  is  much 
more  soluble,  and  crystallizes  with  difficulty.  As  the  alum  employed  in 
dyeing  must  contain  no  iron,  we  understand  why  this  salt  is  not  applica- 
ble. At  present  the  alum  is  being  more  and  more  supplanted  by  alu- 
minium sulphate  and  sodium  aluminate  in  all  practical  operations,  because 
these  chemicals  can  be  procured  perfectly  free  from  iron. 

Aluminium  Phosphate,  AlPO^-j- 4H2O,  is  thrown  out  of  alumin- 
ium salt  solutions  by  sodium  phosphate,  as  a white  gelatinous  precipi- 
tate; this  is  readily  soluble  in  acids,  acetic  excepted.  Turquoise  is  a 
basic  phosphate  of  aluminium  colored  blue  by  a copper  compound. 

Aluminium  Silicates. — The  most  important  of  the  aluminium 
double  silicates,  so  widely  distributed  in  nature,  are  : leucite,  KAl(Si03)2, 
albite  or  soda  feldspar, NaAlSi30g,  ordinary  feldspar — orthoclase,  KAlSi30g, 
and  the  various  micas,  which,  with  quartz,  compose  granite.  When 
these  disintegrate  under  the  influence  of  water  and  the  carbon  dioxide 
of  the  air,  alkaline  silicates  are  dissolved  and  carried  away  by  water, 
while  the  insoluble  aluminium  silicate,  clay,  remains.  Perfectly  pure 
clay  is  white,  and  is  called  kaolin,  or  porcelain  clay  ; its  composition 
mostly  corresponds  to  the  formula,  Al/SiOJg  -f  2H2O  or  Al2Si207  -}- 
2H2O.  When  clay  is  mixed  with  water  a tough,  kneadable  mass  is 
obtained.  By  drying  and  burning,  it  becomes  compact  and  hard,  and  is 
the  more  fire-proof,  the  purer  the  clay.  On  this  depends  the  use  of  clay 
for  the  manufacture  of  earthenware,  from  the  red  brick  to  porcelain. 

30 


354 


INOkCJANIC  CHEMIS'IRY. 


"I'o  produce  porcelain  ii  very  ('me  mixture  of  kaolin,  feldspar  and  quartz  is  cinj)loyed. 
On  strong  ignition,  the  feldspar  fuses,  (ills  the  pores  of  the  clay  and  thus  furnishes  a fused 
transparent  mass — porcelain.  When  it  is  not  so  strotigly  ignited,  it  remains  porous  — 
faience — serving  for  liner  clay  vessels.  'I'o  render  these  impervious  to  water,  they  are 
covered  wdth  glazing.  'I'liis  consists  of  various  readily  fusible  silicates.  Rough  earthen- 
ware vessels  are  constructed  from  impure  clay,  and  they  are  u.sually  glazed  by  throwing 
salt  into  the  ovens  at  the  time  of  burning,  d'he  salt  vaj)orizes  and  forms  an  easily 
fusible  silicate  on  the  surface  of  the  clay,  hydrochloric  acid  being  liberated. 

Ultra7narine. — Tlie  rare  mineral  Lapis  lazuli,  which  was  formerly  em- 
l)loyed  as  a very  valuable  blue  color  under  the  name  of  Ultrainai'ine,  is 
a compound  of  altiminitim  sodium  silicate  with  sodium  j)olysuli)hides. 
At  present  ultramarine  is  prepared  artificially,  in  large  quantities,  by 
heating  a mixture  of  clay,  dry  soda  (or  sodium  stilj)hate),  sulphur  and 
charcoal,  away  from  air.  Green  ultramarine  is  the  product.  This  is 
then  washed  with  water,  dried,  mixed  with  powdered  sulphur  and  gently 
heated  with  air  contact  until  the  desired  blue  color  has  appeared — blue 
ultramarine.  The  chemical  constitution  of  this  blue  coloring  material 
has  not  yet  been  explained.  On  pouring  hydrochloric  acid  over  the  blue 
product,  the  color  disapi)ears  with  liberation  of  sulj)hur  and  hydrogen 
sulphide — this  would  point  to  the  existence  of  a polysulphide.  Violet 
and  red  ultramarines  are  prepared  at  present  by  conducting  dry  hydrogen 
chloride  gas  and  air  over  common  ultramarine  at  100-150°. 


RARE  EARTH  METALS. 

In  some  very  rare  minerals,  like  cerite,  gadolinite,  euxenite,  orthite,  samanskite,  thorite, 
and  monazite,  occurring  principally  in  Sweden,  Norway,  Greenland,  and  the  United 
States,  is  found  a series  of  metals  which,  in  their  entire  deportment,  closely  resemble 
aluminium  (p.  345).  The.se  are  scandium,  yttrium,  cerium,  lanthanum,  neodymium, 
praseodymium,  samarium,  ytterbium,  and  the  more  recent  erbium,  terbium,  gadolinium, 
decipium,  thulium,  holmium  and  dysprosium.  These  generally  form  difficultly  soluble 
oxalates,  and.  are,  therefore,  precipitated  from  solution  by  oxalic  acid.  The  anhydrous 
sulphates  are  quickly  and  easily  soluble  in  cold  water  ; in  warm  water  they  are  slowly  and 
difficultly  soluble.  They  form  double  salts  with  the  alkali  sulphates  ; the  potassium 
double  salts  are  constituted  according  to  the  formula  1^02(804)3. 3X280.^.  Their  solu- 
bility in  water  differs  ; some  of  them  are  insoluble  in  saturated  solutions  of  potassium 
sulphate.  The  formates  are  partly  sparingly  soluble  and  in  part  very  soluble.  The 
hydrous  nitrates  and  chlorides  are  very  soluble  in  water  and  in  alcohol.  The  nitrates  of 
these  metals  are  decomposed  at  different  temperatures  ; this  property  affords  an  excellent 
means  for  their  isolation  and  separation.  Fractional  precipitation  with  ammonia  is  also 
used  for  this  purpose.  The  salts  and  solutions  of  the  individual  metals  are  in  part 
colorless  and  ])artly  colored.  Most  of  the  latter  show  absorption  spectra.  The  position 
of  the  ab.sorjilion  lines  in  the  spectrum  — their  wave-length  — is  characteristic  for  that 
particular  element  and  is  a sure  and  excellent  means  for  its  recognition.  The  oxides  of 
many  of  these  metals  when  heated  in  the  Bunsen  flame  diffu.se  great  quantities  of  light, 
d'he  glowing  oxides  of  .some  of  the  elements  whose  .salt  .solutions  show  absorption  spectra 
give  a di.scontinuous  spectrum  with  bright  bands  whose  positions  corre.spond  exactly  to  the 
ab.sorptioti  lines. 

8candium,  yttrium,  lanthanum,  cerium,  samarium,  and  ytterbium  have  been  most  accu- 
rately investigated.  'Fheir  atomic  weights  are  very  approximately  correct.  The  most  in- 
teresting of  the  grouj)  is  scandium,  atomic  weight  44.1.  It  fills  out  the  gap  between 
calcium  and  titanium.  It  coincides  in  all  its  properties  with  those  deduced  theoretically 
from  the  periodic  system  by  Mendelejeff  for  the  element  ekaboron  (compare  Gallium). 


RARE  EARTH  METALS. 


355 


Scandium,  Sc  = 44.1,  discovered  by  Nilson  in  1879,  is  contained  in  euxenite  and 
gadolinite  ; it  has  not  yet  been  obtained  in  a free  condition.  Its  oxide,  Sc^Oy,  is  obtained 
by  igniting  the  hydroxide  or  nitrate,  and  is  a white,  infusible  powder  (like  magnesia  and 
oxide  of  beryllium).  Its  specihc  gravity  equals  3.86.  The  hydroxide,  Sc(OM).j,  is 
precipitated  as  a gelatinous  mass  from  its  .salts  by  the  alkalies,  and  is  insoluble  in  an 
excess  of  the  latter.  The  nitrate  crystallizes  in  little  prisms,  and  is  easily  decomposed  by 
heat.  The  potassium  double  sulphate,  Sc.^(SO^)3.  3K2SO4,  is  soluble  in  warm  water,  but 
not  in  a solution  of  potassium  sulphate.  The  chloride  affords  a characteristic  spark  spectrum. 

Yttrium — Y = 89,  discovered  by  Mosander  in  1843  in  the  yttria  of  gadolinite,  has 
been  obtained  by  electrolyzing  the  anhydrous  fused  chloride.  It  is  a gray  powder  which 
decomposes  water  in  the  cold.  Cleve  and  Hoglund  first  obtained  many  of  its  compounds 
in  a pure  state.  It  occurs  principally  in  the  gadolinite  from  ytterby  (about  35  per  cent. 
Y2O3).  Its  oxide  is  a white,  infusible,  strong  base,  which  is  capable  of  expelling  ammonia 
from  its  salts.  Its  specific  gravity  is  5.04.  Its potasshan  double  sulphate,  Y2(S()^)3. 3X280^, 
is  soluble  in  a potassium  sulphate  solution,  and  in  this  manner  it  can  be  readily  separated 
from  cerium,  lanthanum,  didymium  and  samarium.  \X.^nitrate,  Y(N03)3  “h  forms 

large,  deliquescent  crystals.  The  chloride,  YCI3  -j-  6H2O,  forms  large  prisms,  and  gives 
a brilliant  spark  spectrum  rich  in  lines. 

Lanthanum,  La  = 138  (from  ?iav6dveiv,  to  be  hidden),  was  discovered  in  cerite  by 
Mosander  (1839).  It  can  be  obtained  by  the  electrolysis  of  its  fused  chloride.  It  resem- 
bles iron  in  color  and  in  luster.  It  oxidizes  in  the  air  and  in  a flame  burns  with  a bright 
light.  Specific  gravity  = 6. 16.  The  oxide  La203  is  brilliant  white  ; it  combines  directly 
with  water  to  the  hydrate  lLa.{OH).^,  which  reacts  alkaline  and  absorbs  carbon  dioxide  from 
the  air.  It  is  a powerful  base  which  expels  ammonia  from  ammonium  salts  in  the  cold. 
The  anhydrous  sulphate,  La2( 80^)3,  dissolves  readily  in  water  of  4°.  When  its  saturated 
solution  is  heated  to  40°  most  of  it  separates  as  the  hydrous  salt.  The  following  sul- 
phates are  known  : La2(804)3  -\-  6H2O  and  La2(804)3  -f  9H2O.  The  potassium  double 
salt  is  sparingly  soluble  in  water  and  insoluble  in  a saturated  potassium  sulphate  solution, 
La2(804)3.  31^2^04.  The  nitrate,  La(N03)3  • 6H2O,  consists  of  large  prismatic  crystals, 
which  lose  water  on  heating.  The  salt  melts  when  more  strongly  heated,  and  on  cooling 
becomes  like  dust.  It  decomposes  at  high  temperatures.  The  anhydrous  chloride, 
LaCl3,  is  a crystalline  mass,  which  deliquesces  in  moist  air,  forming  crystals  of  the  hydrous 
chloride,  2LaCl3  . 15H.2O.  The  chlo7'ide  shows  a very  bright  spark  spectrum  with  numer- 
ous lines.  Pure  lanthanum  oxide  exposed  in  a vacuum  to  kathode  rays  shines  with  a 
powerful  light  and  before  the  spectroscope  shows  a brilliant  band  spectrum.  The  lan- 
thanum salts  are  colorless  ; they  do  not  have  an  absorption  spectrum. 

Cerium,  Ce  — 140,  discovered  simultaneously  in  1803  by  Klaproth  and  Berzelius 
and  named,  as  was  also  the  mineral,  from  the  newly  discovered  planet  Ceres,  occurs  in 
cerite  (60  per  cent.),  and  is  also  obtained  by  the  electrolysis  of  the  chloride.  It  is  very 
similar  to  lanthanum,  but  at  ordinary  temperatures  is  more  stable  in  the  air  than  the  latter  ; 
it  burns  much  more  readily,  so  that  .broken -off  particles  of  it  inflame  of  their  own  accord. 
The  specific  gravity  of  the  fused  metal  is  6.72.  Besides  the  salts  of  the  sesquioxide, 
it  forms  some  from  the  dioxide,  The  former  in  which  cerium  is  trivalent  are  colorless, 

while  the  latter — the  ceric — are  colored  yellow  or  brown  ; red  ceric  hydroxide,  Ce(OH)4, 
is  precipitated  from  the  former,  on  the  addition  of  hypochlorites.  A little  aqueous  hydro- 
fluoric acid  will  convert  the  ceric  hydroxide  into  cej-ium  tet)'aJluo7'ide,Ce¥\^  T H2O.  The 
tetrachloride,  CeCl4,  is  also  known  in  double  salts.  These  compounds  indicate  that  cerium 
is  quadrivalent  and  that  it  probably  belongs  to  the  fourth  group  of  the  periodic  system 
(p.  246). 

Didymium,  Di  — 142,  was  found  in  cerite  by  Mosander,  and  was  regarded  as  an 
element  until  1885.  Auer  von  Welsbach  then  succeeded  in  resolving  it  into  two  new 
elements,  which  he  called  Neodymium  = Nd  (144?)  and  Praseodymium  — Pr 
(140?).  The  oxide  of  the  old  didymium  (from  6i6v/w/,  twins,  a constant  associate  of 
cerium  and  lanthatuim)  has  a bright  gray  color  after  strong  ignition,  that  of  neodymium 
a bluish  color  and  that  of  praseodymium  a greenish-white  color.  While  the  salts  of  old 
didymium  had  an  amethyst-red  color,  those  of  neodymium  are  a rose- violet  and  those  of 
praseodymium  are  an  intense  leek-green.  'I'he  salts  and  .solutions  of  the  old  didymium 
.showed  a magnificent  absorption  spectrum  and  its  components  share  this  with  it.  'I'he 
same  is  true  of  the  spark  spectra.  'I'he  most  important  derivatives  of  praseodymium 


INORGANIC  CllKMISTRY. 


were  recenlly  prepared  and  studied  by  von  Scheie.  I(  lias  two  oxides:  I’rO,^  aiul  I’r,/ 

Its  salts  follow  the  type  I’rX,  ; the  praseodymium  in  them  is  pronouncedly  Irivalent  (Z.  f. 
anorfr.  Ch.  (1898)  17,  310  ; 18,  352). 

Samarium,  Sm  150,  was  discovered  by  M.  1 )elafontaine  and  almost  simultaneously 
by  Lecoq  de  Hoisbaudran  in  1878  on  studying  unknown  absorption  bands  given  by  the 
samarskile  Irom  North  Carolina.  It  has  also  been  found  in  cerite,  gadolinite  and  orthite. 
Marignac  (1880)  prepared  the  oxide  and  some  salts  in  a pure  state.  1’.  'I'.  Cleve  has  made 
an  exhaustive  investigation  of  it.  d'he  metal  has  not  yet  been  prepared.  'I'he  oxide  is 
pure  white  and  its  specific  gravity  is  8.38.  In  the  lUmsen  flame  it  shows  a remarkable 
light-emission  jiower  and  a discontinuous  .spectrum  with  bright  bands.  Samarium  .salts 
are  light  yellow  in  color.  'I'hey  and  their  solutions  exhibit  a characteristic  absorption 
spectrum.  'The  sulpItaiCy  Sm.^(SOj).j  | 811.^0,  consists  of  small  crystals,  dis.solving  with 
difficulty  in  water.  Its  double  .salt  with  potassium  sulphate  is  sparingly  soluble  in  water 
and  almost  insoluble  in  a saturated  potassium  sulphate  .solution,  d'he  ni/rafr,  .Sm(N()3).„ 
crystallizes,  melts  easily  and  is  decomposed,  d'he  chloride,  SmClj  611./),  forms  large 
tabular  crystals  ; its  .S])ark  spectrum  is  very  distinct. 

Ytterbium,  Vb  — 173.  Marignac  (1878)  found  that  its  oxide  was  the  chief  constituent 
of  erbium  oxide  which  Hahr  aiul  Bunsen  had  regarded  as  a distinct  element.  It  is  ob- 
tained by  the  repeated  incomplete  decomposition  of  a mixture  of  the  nitrates  of  these 
earths  with  careful  heating.  Ytterbium  and  scandium  oxides  being  the  feeblest  ba.ses, 
are  first  eliminated  and  can  then  be  se])arated  by  means  of  potassium  sulphate.  Ytterbium 
oxide,  Yb./).,,  is  a white,  infusible  ])owder  of  specific  gravity  9.17.  It  is  readily  soluble 
in  acids;  its  salts  are  color]e.ss  and  do  not  show  an  absorption  spectrum.  Ytlerlduni 
sulphate,  Yb.^(.S()^)3  811^0,  forms  large  pri.sms,  stable  in  the  air.  Its  double  salt  with 

potassium  sulphate  is  readily  .soluble  in  water  and  in  a .saturated  potassium  sulphate 
solution.  Ytterbium  chloride  gives  a spark  spectrum  rich  in  line.s. 

Erbium  — Er.  Mosander  (1843)  demonstrated  that  the  supposedly  simple  yttria  was 
in  fact  a mixture  of  the  three  earths — erbia,  terbia  and  ytterbia.  The  erbia  of  to-day 
is  the  rose-red  oxide  from  the  old  yttria,  which  Mosander  called  terbia,  and  which  Bahr 
and  Bunsen  ( 1866),  as  well  as  Cleve  and  Hoglund  (1873)  prepared  in  large  quantities 
from  gadolinite  and  regarded  it  as  elementary.  However,  Marignac  in  1878  obtained 
ytterbium  from  it  and  Nilson  in  1879  scandium.  The  investigations  continued  in  that 
line  proved  that  it  could  be  still  further  simplified.  The  old  erbia  gave  an  ab.sorption 
spectrum  very  rich  in  lines  and  the  fact  that  on  partly  decomposing  its  nitrate  by  fusion 
or  by  the  fractional  precipitation  of  its  nitrate  .solution  with  dilute  ammonia,  fractions  were 
obtained  in  which  the  ab.sorption  bands  appeared  in  different  number  and  strength, 
induced  P.  T.  Cleve  to  assume  in  it  the  presence  of  three  elements — the  real  Erbium  = 
Er  (166),  Thulium  = Tm  and  Holmium  ^IIo.  Then  Lecoq  de  Boisbaudran  by  similar 
observations  discovered  that  holmium  also  contained  Dysprosium  = Dy  {SvaTTpoatro^, 
difficult  to  reach).  Nilson  and  Kriiss  regard  erbium,  thulium,  holmium,  and  dysprosium 
as  capable  of  further  division. 

Terbium  = Tr.  The  oxide  (terbia)  was  discovered  by  Mosander  in  the  old  yttria, 
but  called  erbia  by  him  at  that  time.  It  occurs  in  small  quantities  in  the  gadolinite  of 
Ytterby  in  Sweden.  The  names,  yttrium,  erbium,  terbium,  and  ytterbium,  are  transposi- 
tions of  the  syllables  of  the  word  Ytterby.  The  samarskite  from  North  Carolina  contains 
great  quantities  of  terbia.  The  metal  has  not  been  prepared.  Terbia  is  a dark  orange-red 
powder.  It  dis.solves  in  hydrochloric  acid  with  evolution  of  chlorine.  It  loses  its  color 
when  heated  in  hydrogen,  also  suffers  a slight  loss  in  weight  and  becomes  white.  Very 
probably  the  orange-colored  oxide  is  a peroxide.  The  terbium  salts  are  colorless  ; they 
have  no  absorption  spectrum.  Its  formate  dissolves  with  difficulty  in  water.  The  double 
sulphates  with  potassium  and  sodium  sulphate  are  .s[)aringiy  soluble  in  saturated  potassium 
sulidiate  and  sodium  sulphate  solutions.  The  chloride  does  not  yield  a spark  spectrum, 
d'erbia  is  ])robably  identical  with  the  oxide  of  philippium,  di.scovered  by  M.  Delafontaine, 
and  the  oxide  of  mosandrium,  by  |.  Lawrence  Smith.  Lecoq  de  Boisbaudran  considers 
that  terbia  contains  two  additional  earths,  which  he  designates  temporarily  7m  and  7,3. 

Gadolinium,  Cd  ^ 156?  Marignac  discovered  its  oxide  in  the  North  Carolina  samar.s- 
kite.  It  is  also  |)resenl  in  so.me  ortliites.  It  is  an  almost  white  ]-)owder,  very  .soluble  in 
acids.  'The  salts  and  salt  solutions  do  not  exhibit  an  absorption  spectrum  ; the  chloride 
(l(jes  not  give  a spark  spectrum.  Expo.sed  to  the  kathode  rays  of  an  induction  machine 


THE  GALLIUM  GROUP.  357 

in  an  almost  vacuous  tube  gadolinia  gives  out  a fiery  red  color,  and  in  the  spectroscope 
has  a spectrum  with  a bright  orange-red  line,  whose  wave  length  A = 6094. 

Decipium,  Dp  =;  lyi  ? Its  oxide  decipia  (from  decipere,  to  deceive)  was  discovered 
in  the  samarskite  of  North  Carolina  by  Delafontaine.  It  is  colorless  ; its  salts  also  are 
colorless  and  show  no  absorption  spectrum.  Its  double  sulphate  with  potassium  sulphate 
is  insoluble  in  a saturated  potassium  sulphate  solution.  Decipium  does  not  give  a spark 
spectrum. 


2.  THE  GALLIUM  GROUP. 


The  three  heavy  metals,  gallium,  indium,  and  thallium,  bear  the  same 
relation  to  aluminium  which  copper,  silver,  and  gold  bear  to  sodium 
and  zinc  cadmium,  and  mercury  to  magnesium. 


Cu 

63.6 

Zn 

65.4 

Ga 

70 

Ge 

72 

As 

75 

Ag 

107-93 

Cd 

112 

In 

114 

Sn 

118.5 

Sb 

120 

Au 

197.2 

Hg 

200.3 

Tl 

204. 1 

Pb 

206.9 

Bi 

208.5 

They  constitute  the  corresponding  members  of  the  three  great  periods, 
and  as  a second  sub-group  attach  themselves  to  aluminium,  while  cerite 
metals  form  the  first,  more  basic  sub-group  (p.  354).  The  entire  character 
of  the  three  elements  under  consideration  is  influenced  by  this  position 
in  the  periodic  system,  because  regular  relations  appear  in  all  directions, 
as  may  be  observed,  for  example,  in  the  specific  gravities,  melting  points, 
and  other  physical  properties  in  the  free  metals  : 


Ga 

In 

Tl 

Atomic  weight, 

70 

II4 

204.1 

Specific  gravity, 

5-9 

7-4 

II. 8 

Melting  point,  

30° 

176° 

290° 

Being  members  of  group  III  of  the  periodic  system,  Ga,  In,  and  T1 
yield  compounds  of  the  trivalent  form,  and  these  are  analogous  to  those 
of  aluminium  in  many  respects. 

Thallium,  like  other  elements  with  high  atomic  weights  (Au,  Hg,  Pb), 
exhibits  great  variations  from  the  group  properties  (p.  321).  It  yields, 
for  example,  not  only  derivatives  of  the  form  TIX3,  but  also  those  of 
'FIX.  If  we  include  thallium  as  a member  of  the  last  great  period  (Pt, 
Au,  Hg,  Tl,  Pb,  Bi),  we  will  discover  that,  as  in  the  case  of  the  other 
metals  of  this  series,  a remarkable  regularity  underlies  all  its  forms  of 
combination — the  highest  as  well  as  the  lowest : 

PtCb  AuCl  IlgCl  TlCl  PbCb  BiCla 

PtCl,  AUCI3  IlgCl,  TICI3  PbCl,  BiXg. 


358 


INORGANIC  CHEMISTRY. 


1.  GALLIUM. 

Ga  = 70. 

Gallium  was  discovered  iii  zinc-blende  from  I’ierrefitte,  in  1875,  by  Lecoq  de  lioisbaud- 
ran,  by  means  of  the  spectrosc()i)e.  Since  then  it  has  been  observed  in  other  blendes, 
and  in  the  clay  iron  ore  of  Yorkshire,  as  well  as  in  the  pig-iron  of  Middlesborough.  It 
is  identical  with  Linnemann’s  austriuin.  As  early  as  the  year  1869,  Mendelejeff,  taking 
the  table  of  the  periodic  system  devised  by  him  as  basis,  predicted  tlie  existence  of  a metal 
(standing  between  aluminium  and  indium,  with  an  atomic  weight  of  nearly  69),  which  he 
named  Kka-aluminimn.  Its  i)roperties  were  necessarily  deduced  from  its  position  in  the 
periodic  system.  All  the  jnoperties  of  gallium  known  at  that  time  agreed  with  those  of 
eka-aluminium,  and  it  seemed  very  j^robable  that  this  element,  which  had  been  theoreti- 
cally established,  was  in  reality  gallium. 

As  yet  gallium  has  only  been  found  in  very  small  (juantity,  and  is  but  imperfectly 
investigated.  It  is  characterized  by  a spectrum  consisting  of  two  violet  lines.  Separated 
by  electrolysis  from  an  ammotiiacal  solution  of  its  sulphate,  it  is  a white,  hard  metal,  of 
specific  gravity  5.9,  with  a fusing  point  of  4-30°.  It  is  only  .superficially  oxidized  in  the 
air,  not  altered  by  water,  and  is  not  volatile  up  to  a red  heat.  Like  aluminium,  it  is 
scarcely  attacked  by  nitric  acid,  but  dis.solves  readily  in  hydrochloric  acid,  as  well  as  in 
caustic  pota.sh  and  ammonia  with  the  evolution  of  hydrogen. 

Gallium  and  aluminium  unite  to  form  alloys  ; those  rich  in  the  former  metal  are  liquid 
at  the  ordinary  temperature,  and  decompo.se  water  as  energetically  as  metallic  .sodium. 

Gallium  Oxide,  GajOg,  is  obtained  by  igniting  the  nitrate.  It  is  a white  mass  which 
sublimes  when  heated  in  a current  of  hydrogen.  'I'he  hydroxide,  Ga(OII)3,  is  thrown 
out  of  the  solutions  of  its  salts  by  the  alkalies  as  a white  flocculent  precipitate,  readily 
soluble  in  an  excess  of  the  precipitant,  but  rather  difficultly  soluble  in  ammonium 
hydroxide. 

Gallium  Chloride,  GaCIg,  is  produced  on  heating  gallium  in  a current  of  chlorine 
gas  or  in  hydrogen  chloride  ; it  forms  large  colorless  crystals  which  fu.se  at  75°,  sublime 
at  about  60°  and  boil  at  215-220°.  Its  vapor  density  above  440°  corresponds  to  the 
formula  GaCl^,  at  270°  very  closely  to  Ga.^Clg.  The  chloride  fumes  and  deliquesces  in  the 
air,  like  aluminium  chloride,  and  decomposes  in  the  evaporation  of  its  aqueous  solution. 

Gallium  Nitrate,  Ga(N03)3,  and  Gallium  Sulphate,  Ga.3(S04)3,  are  crystalline  and 
very  deliquescent.  The  latter  forms  a double  salt  with  ammonium  sulphate — similar  to 
the  alums  : 

Ga3(SOj3.(NH,)3SO,  + 24H.P. 

Hydrogen  sulphide  only  precipitates  gallium  from  acetic  acid  solutions. 

Gallium  Dichloride,  GaCl2,  is  also  known.  It  results  on  heating  the  trichloride 
with  metallic  gallium.  It  consists  of  colorless  crystals,  melts  at  164°  and  distils  at  535°- 
Its  vapor  density  at  1000°  corresponds  to  the  formula  GaClj. 


2.  INDIUM. 

In  = 114. 

Owing  to  its  resemblance  to  zinc,  indium  was  regarded  as  a bivalent  metal,  and  its 
compounds  were  suj)po.sed  to  have  the  formula  InX.^ ; this  fixed  its  atomic  weight  at  75-6. 
'file  specific  heat,  however,  made  the  atomic  weight  one  and  a half  times  as  large  (p. 
251).  Hence  it  is  trivalent  and  its  derivatives  are  constituted  according  to  the  form  10X3. 
It  belongs  to  the  groiq)  of  aluminium,  and,  in  its  derivatives,  manifests  some  similarity  to 
this  metal. 

It  was  discovered,  in  1863,  by  Reich  and  Richter,  by  the  aid  of  spectrum  analysis. 
Its  S[)ectrum  is  characterized  l)y  a very  bright  indigo-blue  line,  hence  its  name.  It  only 
occurs  in  very  minute  ([uantilies  in  .some  zinc-blendes  from  Freiberg  and  the  Hartz. 

It  is  a silver-white,  soft  and  tenacious  metal,  of  s])(f(:ific  gravity  ']  \2.  It  melts  at 
176°  and  distils  at  a wliin?  heat.  At  ordinary  temperatures  it  is  not  altered  in  the  air; 


THALLIUM.  359 

heated,  it  burns  with  a blue  flame  to  indium  oxide.  It  is  difficultly  soluble  in  hydro- 
chloric and  sulphuric  acids,  but  dissolves  readily  in  nitric  acid. 

Indium  Chloride,  InClg,  results  from  the  action  of  chlorine  on  metallic  indium,  or 
upon  a heated  mixture  of  indium  oxide  and  carbon.  It  sublimes  in  white,  shining 
leaflets,  which  deliquesce  in  the  air.  Its  vapor  density  at  600-800°  corresponds  to  the 
formula  InClg,  but  above  840°  it  undergoes  a gradual  decomposition.  It  does  not 
decompose  when  its  aqueous  solution  is  evaporated. 

Indium  Oxide,  In^Og,  is  a yellow  powder  resulting  from  the  ignition  of  the 
hydroxide. 

Indium  Hydroxide,  In(OH)g,  is  precipitated  as  a gelatinous  mass,  by  alkalies,  from 
indium  solutions.  It  is  soluble  in  sodium  and  potassium  hydroxides. 

Indium  Nitrate,  In(NOg)g,  crystallizes  with  three  molecules  of  water,  in  white 
deliquescent  needles.  *, 

Indium  Sulphate,  In2(S04)g,  remains  on  evaporating  a solution  of  indium  in  .sul- 
phuric acid  as  a gelatinous  mass,  with  three  molecules  of  water.  It  forms  an  alum  w'ith 
ammonium  sulphate. 

Indium  Sulphide,  lugSg,  is  precipitated  by  hydrogen  sulphide  as  a yellow-colored 
compound  from  indium  solutions.  It  combines  to  sulpho-salts,  InS . SK  (from 

InO.OH)  with  the  sulphides  of  the  alkalies. 

Indium  also  forms  a dichloride  and  monochloride. 

Indium  Dichloride,  InClg,  is  produced  when  metallic  indium  is  heated  in  a current 
of  hydrogen  chloride.  It  is  a white  crystalline  mass,  which  on  exposure  to  a more 
intense  heat  becomes  a yellow  liquid  and  sublimes.  Its  vapor  density  at  1000-1400° 
corresponds  to  the  formula  InClg.  Water  decomposes  it  at  once  into  indium  trichloride 
and  metallic  indium  : sInClg  = 2lnCl3  + In. 

Indium  Monochloride,  InCl,  like  gallium  dichloride  (p.  358),  results  when  the 
dichloride  is  heated  with  metallic  indium.  It  is  a crystalline,  reddish-yellow  mass,  which 
has  a reddish-black  color  when  fused.  Its  vapor  density  at  1100-1400°  corresponds  to 
the  formula  InCl,  Water  decomposes  it  into  the  trichloride  and  metallic  indium : 
3lnCl  = InClg  -)-  2ln. 


3.  THALLIUM. 

T1  = 204.1. 

Thallium  is  rather  widely  distributed  in  nature,  but  in  very  small  quantity.  The  very 
rare  mineral  crookesite  contains  17  per  cent,  of  the  metal,  together  with  copper,  selenium 
and  silver.  It  is  often  found  with  potassium  in  sylvite  and  carnallite,  in  mineral  springs, 
and  in  some  varieties  of  pyrite  and  zinc-blendes.  When  these  pyrites  are  roasted  for  the 
production  of  sulphuric  acid,  according  to  the  chamber  process,  the  thallium  deposits  as 
soot  in  the  chimney  and  in  the  chamber  sludge,  and  was  discovered  in  the  latter,  almost 
simultaneously,  by  Crookes  (1861)  and  Lamy  (1862  ),  by  means  of  the  spectroscope. 

To  get  the  thallium,  the  chimney-dust  is  boiled  with  water  or  sulphuric  acid,  and 
thallous  chloride  precipitated  from  the  solution  by  hydrochloric  acid.  The  chloride  is 
then  converted  into  sulphate,  and  the  metal  separated  from  the  latter  by  means  of  zinc  or 
the  electric  current. 

Thallium  is  a white  metal,  as  .soft  as  sodium,  and  has  the  specific  gravity  11.8.  It 
melts  at  290°,  and  distils  at  a white  heat.  It  oxidizes  very  rapidly  in  moist  air.  It 
does  not  decompose  water  at  ordinary  temperatures.  It  is,  therefore,  best  preserved 
under  water  in  a clo.sed  vessel.  By  air  access  it  gradually  di.ssolves  in  the  water,  forming 
thallous  hydroxide  and  carbonate.  Heated  in  the  air  it  burns  with  a beautiful  green 
flame  who.se  spectrum  shows  a veiy  inten.se  green  line,  hence  the  name  thallium,  from 
i^a^Aor,  green.  Thallium  dissolves  readily  in  sulphuric  and  nitric  acids,  but  is  only 
.slightly  attacked  by  hydrochloric  acid,  owing  to  the  insolubility  of  thallous  chloride, 

I III 

Thallium  forms  two  series  of  com[)ound.s  : thallous,  TlX,  and  thallic,  TlXg.  The 
first  are  very  similar  to  the  compounds  of  the  alkalies  (and  also  to  those  of  silver).  The 
solubility  of  the  hydroxide  and  carbonate  in  water  shows  this  ; their  solutions  have  an 
alkaline  reaction.  Again,  many  thallous  salts  are  i.somorphous  with  those  of  potassium. 


360 


INORGANIC  CIIKMIS’IRV. 


iiiul  form  similar  (loul)le  sails  (see  Alums,  p.  352).  In  the  insoluhilily  of  ils  sulplmr 
and  haloj^en  compounds,  univalent  thallium  api)roaches  silver  and  lead. 

In  its  comi)oundsof  the  form  'I'lXj  thallium  is  trivalent,  like  aluminium,  hut  otherwise 
shows  .scarcely  any  similarity  to  the  latter. 

'rhallium  compounds  are  poi.sonou.s.  'i'hey  are  used  in  making  thallium  glass,  which 
refracts  more  strongly  than  lead  glass. 


THALLOUS  COMPOUNDS. 

Thallous  Oxide,  Tl.^0,  is  formed  by  the  oxidation  of  thallium  in  the  air,  or  by  heat- 
ing the  hydroxide  to  ioo°^  It  is  a black  powder  which  di.ssolves  in  water  with  formation 
of  the  hydroxide. 

Thallous  Hydroxide,  'ri(OII),  may  be  prepared  by  decomposing  thallous  sulphate 
with  an  equivalent  amount  of  barium  hydroxide,  and  crystallizes  with  one  molecule  of 
water  in  yellowish  prisms.  It  di.ssolves  readily  in  water  and  alcohol,  yielding  strong 
alkaline  solutions. 

Thallous  Chloride,  TlCl,  forms  on  heating,  the  metal  in  hydrochloric  acid  gas  and 
is  thrown  down  from  solutions  of  thallous  salts  by  hydrochloric  acid  as  a white,  curdy 
precipitate,  which  is  difficultly  soluble  in  water.  It  .separates  in  .small  crystals  from  the  hot 
solution.  It  fu.ses  at  427°,  and  boils  at  about  715°.  Like  ])ota.ssium  chloride,  it  forms  an 
insoluble  double  salt  with  jdatinic  chloride,  rt(jl^.2'riCl.  'J'/iallous  hrofnide  forms  a white, 
and  thallous  iodide  a yellow  ]u-eci{)itate.  'I'he  latter  is  not  .soluble  in  potassium  iodide. 

Thallous  Sulphate,  'Fl.^SO^,  crystallizes  in  rhombic  pri.sm.s,  isomorphous  with  potas- 
sium sulj)hate.  It  dissolves  in  20  parts  of  water  at  ordinary  temperatures.  It  affords 
double  salts  with  the  sulphates  of  the  metals  of  the  magnesium  group,  of  ferrous  oxide, 
of  cupric  oxide,  etc.  (p.  315),  e.  g.,  MgSO^.  Tl^SO^  + 6H,^0  ; these  are  perfectly  similar 
and  analogous  to  the  corresponding  double  salts  of  potassium  and  ammonium.  It 
forms  thallium  alum  with  the  .sulphates  of  the  .sesquioxides  of  the  iron  group,  e.  g., 
T1A1(S04)2  -f-  12H2O  ; these  are  similar  to  potassium  alum,  KA1( 80^)2  + I2li20. 

Thallous  Carbonate,  Tl^COg,  is  obtained  from  the  oxide  by  the  absorption  of  carbon 
dioxide  ; it  crystallizes  in  needles,  which  dissolve  at  ordinary  temperatures  in  20  parts  of 
water.  The  solution  has  an  alkaline  reaction. 

Thallous  Sulphide,  Tl^S,  is  precipitated  from  thallous  salts  by  hydrogen  sulphide 
as  a black  compound,  insoluble  in  water. 


THALLIC  COMPOUNDS. 

Thallic  Chloride,  TICI3,  is  produced  by  the  action  of  chlorine  upon  thallium  or 
thallous  chloride  in  water,  and  is  very  .soluble  in  the  latter.  It  decomposes  at  100°  into 
thallous  chloride  and  chlorine.  The  alkalies  precipitate  from  its  solutions  thallic 
hydroxide,  'riO.OII,  a brown  powder,  which,  at  100°,  passes  into  thallic  oxide, 
Tl./  >3.  Further  heating  decomposes  the  latter  into  thallous  oxide  and  oxygen. 

'The  oxide  and  hydroxide  are  soluble  in  hydrochloric,  nitric,  and  sulphuric  acids, 
forming  'I'lCl.^,  Tl(NO.j).,, 

On  conducting  chlorine  through  a solution  of  thallic  hydroxide  in  jiotassium  hydroxide, 
it  assumes  an  inten.se  violet  color,  due  ju’obably  to  the  formation  of  the  potassium  salt  of 
thallic  acid,  the  compo.sition  of  which  is  yet  unknown. 


METALS  OF  THE  FOURTH  GROUT. 


361 


METALS  OF  THE  FOURTH  GROUP. 

The  elements  of  group  IV  in  the  periodic  system  (p.  246), 

Ti  = 48. 1 Zr  = 90.6  Ce  = 140  Th  = 282 

C = 12.00  Si  = 28.4 

Ge  = 72  Sn  =118.5  Pb  = 206.9, 

show  the  same  analogies  that  were  observed  with  the  members  of  group 
III  (p.  345).  Their  character  is,  however,  non-metallic ; their  deri\  a- 
tives  are  chiefly  of  the  types  MeX^  and  Me02,  of  which  the  latter  a;e 
acid  (p.  258). 

The  first  two  elements,  carbon  and  silicon,  with  low  atomic  weights, 
belong  to  the  two  small  periods  and  are  true  metalloids.  Their  oxides  and 
hydroxides  are  acid  in  nature.  The  first  more  basic  sub-group  comprises 
titanium,  zirconium,  cerium,  and  thorium.  They  constitute  the  fourth 
members  of  the  large  periods.  Their  compounds  are  almost  exclusively 
of  the  type  Me02,  similar  to  the  silicon  derivatives,  and  are 

usually  discussed  (with  the  exception  of  cerium)  with  the  metalloids  after 
silicon  (pp.  238,  355).  The  other  sub-group  consists  of  more  electro- 
negative heavy  metals  : germanium,  tin,  and  lead.  These  constitute  the 
transition  from  the  elements  in  group  III,  corresponding  to  them,  to  those 
of  group  V : 

Ga  70  Ge  72  As  75 

In  1 14  Sn  118.5  Sb  120 

Tl  204.1  Pb  206.9  Pi  208.5. 

Their  intermediate  position  accounts  for  their  metalloidal  character.  In 
this  group,  as  in  all  other  groups,  it  is  noticed  that  as  the  atomic  weight 
rises  (from  germanium  to  lead)  there  is  a successive  rise  in  metallo-basic 
character.  All  three  members  form  dioxides, 

GeOg  SnOj  Pb02, 

which  may  be  viewed  as  anhydrides  of  the  acids 

H2Ge03  H2Sn03  H2Pb03. 

These  are  perfectly  analogous  to  silicic  acid,  but  their  stability  and  acidity 
diminish  as  the  atomic  weights  of  their  basal  elements  increase. 

Lead  dioxide,  Pb02,  combines  with  bases  (especially  the  alkalies),  form- 
ing salts  of  plumbic  acid,  , potassium  plumbate,  K2Pb03.  These  are 
not  very  stable ; water  decomposes  them  into  their  components.  Lead 
dioxide  unites  with  difficulty  with  acids  to  yield  salts.  When  digested  with 
sulphuric  acid  it  liberates  an  atom  of  oxygen,  and  forms  salts  of  lead  mon- 
oxide, PbO.  It  yields  chlorine  with  hydrochloric  acid,  but  in  the  cold 
the  unstable  tetrachloride,  PbCl^,  can  be  obtained.  In  this  respect  lead 
dioxide  resembles  the  peroxides,  mangane.se  peroxide,  MnOj,  and  is 
commonly  known  as  lead  peroxide.  Plowever,  the  salts,  Me2Pb()3,  PbCl^, 
and  the  organo-metallic  compounds,  such  as  Pb(CH3)^,  argue  in  favor  of 
quadrivalent  lead,  and  make  it  perfectly  analogous  to  tin  (p.  258). 

31 


362 


INORGANIC  CHEMISTRY. 


The  elements  of  this  grou})  yield  monoxide  derivatives, 

GeO  SnO  PbO. 

These  are  commonly  known  as  ous  compounds.  They  are  basic  and 
only  form  salts  with  acids,  'bhe  basicity  and  stability  of  their  deriva- 
tives increase  as  the  atomic  weights  rise.  The  german^?//^  and  stann^7;/j 
compounds  are  readily  oxidized  to  derivatives  of  the  dioxide  type,  while 
lead  monoxide,  PbO,  is  a strong  base,  and  forms  very  stable  salts. 


1.  GERMANIUM. 

Ge  = 72. 

This  element  was  discovered  in  1886  by  Cl.  Winkler,  of  Freiberg.  As 
early  as  1871  MendelejeflT,  with  the  periodic  system  as  his  basis,  predicted 
the  existence  of  an  element  with  an  atomic  weight  of  about  73,  which 
corresponded  to  the  then  existing  gap  between  silicon  and  tin;  he 
called  it  ekasilicon  (the  first  analogue  of  silicon).  The  perfect  agree- 
ment of  the  essential  properties  of  germanium  with  those  of  the  theoretical 
ekasilicon  constitutes  a brilliant  confirmation  of  the  law  of  periodicity 
(P-  357)- 

Winkler  discovered  germanium  in  the  very  rare  mineral,  argyrodite.  The  latter  is  a 
double  sulphide  of  germanium  and  silver,  Ge^S  . 3Ag2S.  Penfield  claims  the  formula 
GeS^ . 4Ag,,S  for  argyrodite  as  well  as  for  confieldite.  The  first  is  monoclinic  and  the 
second  isometric.  It  is  also  present  in  minute  quantities  in  euxenite  (together  with 
titanium  and  zirconium)  (Krtiss),  in  samarskite  and  frankeite.  It  maybe  separated  from 
these  minerals  by  fusing  them  with  sulphur  and  soda.  Sodium  sulphogermanate  is  then 
produced,  and  it  is  soluble  in  water  (p.  363). 

To  obtain  free  germanium,  its  dioxide  is  heated  in  a current  of  hydrogen  or  reduced  with 
carbon.  The  product  is  a dark-gray  powder,  which  melts  at  900°,  and  upon  solidifying 
readily  crystallizes  into  beautiful,  grayish-white,  metallic  octahedra.  Its  specific  gravity 
at  20°  equals  5.469.  Its  specific  heat  was  found  equal  to  0.0737  at  100°,  and  at  440°, 
0 0757.  Therefore,  its  atomic  heat  at  100°  is  5.33  and  at  440°,  5.45.  It  increases  very 
slightly  (like  those  of  aluminium  and  silicon)  with  rise  of  temperature  and  is  a little  less 
than  the  mean  atomic  heat  (p.  254). 

Germanium  is  very  stable  in  the  air.  When  ignited  it  burns  with  the  production  of 
white  vapors  of  germanium  dioxide,  GeO.^.  The  metal  (like  silicon)  is  insoluble  in  hydro- 
chloric acid.  Nitric  acid  converts  it  (like  tin)  into  the  hydrate  of  the  dioxide.  It  is  solu- 
ble in  alkalies  upon  fu.sion.  When  heated  in  the  non-luminous  gas  flame,  germanium 
and  its  compounds  do  not  impart  a color  to  the  same.  Its  spectrum  can  only  be  pro- 
duced by  the  action  of  the  induction  .spark. 

Germanium,  like  tin,  forms  derivatives  of  the  oxides  GeO  and  Ge02 ; the  first  are  called 
germanowj  compounds,  the  latter  germanzV,  or  derivatives  of  germanic  acid. 


• GERMANOUS  COMPOUNDS. 

'I'liese  are  not  very  stable,  and  are  readily  oxidized  to  the  higher  form. 

Germanous  Oxide,  GeO,  is  formed  when  the  hydroxide  is  ignited  in  a current  of 
carbon  dioxide.  It  is  a grayish-black  powder.  Germanous  Hydroxide,  Ge(0H)2,  is 
|)rccij)itated  as  a yellow-colored  compound  iq)on  the  addition  of  caustic  alkali  to  the  solu- 
tion of  Ibc  (lichloride.  It  is  soluble  in  hydrochloric  acid. 


TIN.  363 

Germanous  Chloride,  GeClg,  has  not  been  obtained  pure.  It  is  formed  when  hydro- 
chloric acid  gas  acts  upon  heated  germanous  sulphide. 

Germanous  Sulphide,  GeS,  is  a reddish-brown  precipitate  produced  by  the  action  of 
hydrogen  sulphide  upon  the  solution  of  the  dichloride.  It  may  be  obtained  in  grayish- 
black  crystals  by  heating  germanium  sulphide  in  hydrogen  gas.  It  is  soluble  in  hot 
hydrochloric  acid,  forming  the  corresponding  chloride. 


GERMANIC  COMPOUNDS. 

Germanium  Tetrachloride,  GeCl^,  is  formed  by  the  direct  union  of  germanium  with 
an  excess  of  chlorine.  The  metal,  when  gently  heated,  burns  in  an  atmosphere  of  chlo- 
rine, with  a bluish  color.  When  in  powder  form  it  inflames  at  the  ordinary  temperature. 
The  tetrachloride  is  also  produced  if  the  sulphide,  GeSg,  be  heated  together  with  mercuric 
chloride.  It  is  a colorless,  mobile  liquid,  of  specific  gravity  1.887  18°.  It  boils  at 

86°.  It  fumes  strongly  in  moist  air,  and  is  decomposed  by  water  into  hydrochloric  acid 
and  germanic  hydroxide,  Ge(OII)4.  It  is  not  decomposed  by  concentrated  sulphuric 
acid.  Its  vapor  density,  from  300-740°,  corresponds  to  the  molecular  formula,  GeCl^. 

Germanium  Chloroform,  GeHClg,  corresponding  to  ordinary  chloroform,  CHCI3 
(see  p.  159),  is  produced  when  metallic  germanium  is  heated  in  a current  of  hydrochloric 
acid  gas.  It  is  a mobile  liquid,  boiling  at  about  72°.  Its  vapor  density  approximates 
the  molecular  formula  GeHClg.  It  becomes  cloudy  on  exposure  to  the  air,  and  colorless, 
oily  drops  of  Germanium  Oxychloride,  GeOCl2  (?),  separate. 

Germanium  Bromide,  GeBr^,  is  a strongly  fuming  liquid,  which  solidifies  at  0°  to  a 
crystalline  mass. 

Germanium  Iodide,  Gel^,  results  upon  heating  germanium  chloride  with  potassium 
iodide,  or  more  readily  by  conducting  iodine  vapor  over  heated  and  finely  divided  metal. 
It  is  an  orange-colored  solid,  melting  at  144°,  and  boiling  at  400°. 

Germanium  Dioxide,  Ge02,  germanic  anhydride,  is  formed  upon  roasting  the  metal 
or  the  disulphide,  or  by  treating  the  latter  with  nitric  acid.  It  is  a stable,  white  powder, 
of  specific  gravity  4. 70  at  18°.  It  is  slightly  soluble  in  water  (i  part  in  95  parts  at  100°) 
and  imparts  to  the  latter  an  acid  reaction.  Germanic  Hydroxide,  Ge(OH)^,  or 
GeO(OH)2,  Germanic  Acid,  is  produced  by  directly  transposing  the  chloride  with 
water.  It  has  not  been  obtained  perfectly  pure,  as  it  loses  more  or  less  water.  Like 
silicic  acid,  it  is  wholly  acid  in  its  character,  and  only  forms  salts  with  bases.  It  is 
soluble  in  the  hydroxides  and  carbonates  of  the  alkalies,  especially  on  fusion,  while  it  is 
almost  insoluble  in  acids. 

Germanic  Sulphide,  GeS2.  Concentrated  hydrochloric  acid  or  sulphuric  acid  will 
precipitate  it  from  solutions  of  its  sulpho-salts.  It  is  also  formed  when  hydrogen  sulphide 
is  conducted  through  strongly  acidulated  solutions  of  the  oxide.  It  is  a white,  voluminous 
precipitate,  insoluble  in  acids,  but  readily  soluble  in  water.  If  the  precipitate  is  washed 
with  water  it  dissolves.  It  is  reprecipitated  by  acids,  especially  if  hydrogen  sulphide  be 
conducted  through  the  solution.  The  sulphide  dissolves  readily  in  the  fixed  alkaline 
hydroxides  and  ammonia.  It  forms  sulpho-sa\X.s  with  the  alkaline  sulphides.  These  are 
perfectly  analogous  to  the  sulpho-stannates.  Argyrodite  is  an  example  of  this  class, 
Ag^GeSg  + 2Ag2S  (p.  362). 


2.  TIN. 

Sn  = 118.5. 

Tin  {Sfannum)  occMYs  in  nature  principally  as  dioxide  (cassiterite,  tin- 
stone) on  the  Malay  Peninsula,  in  the  islands  of  Banca  and  Bilitong, 
and  in  England  (Cornwall),  Saxony,  India,  and  in  Australia.  To  pre- 
pare the  metal  the  oxide  is  roasted,  lixiviated,  and  heated  in  a furnace 
with  charcoal  : 


SnOj  -f-  2C  — Sn  -p  2CO. 


3^4 


INORGANIC  CHEMISTRY. 


Thus  obtained,  it  usually  contains  iron,  arsenic,  and  other  metals;  to 
purity  it  the  metal  is  fused  at  a low  temperature,  when  the  pure  tin  flows 
away,  leaving  the  other  metals.  The  tin  obtained  in  the  East  Indian 
isles  is  almost  chemically  pure,  while  that  of  England  and  of  Saxony  con- 
tains traces  of  arsenic  and  coj)per. 

Tin  is  an  almost  silver-white,  strongly  lustrous  metal,  with  a specific 
gravity  of  7.3.  It  possesses  a crystalline  structure;  and  when  a rod  of 
it  is  bent  it  emits  a peculiar  sound  (tin  cry),  due  to  the  friction  of  the 
crystals.  Ui)on  etching  a smooth  surface  of  tin  with  hydrochloric  acid, 
its  crystalline  structure  is  recognized  by  the  appearance  of  remarkable 
striations.  At  low  temiieratures  i)erfectly  pure  compact  tin  ])asses  grad- 
ually into  an  aggregate  of  small  quadratic  crystals.  The  metal  is  tolera- 
bly soft,  and  very  ductile,  and  may  be  rolled  out  into  thin  leaves  (tin-foil). 
It  becomes  brittle  at  200°,  and  may  then  be  powdered.  It  fuses  at  231°, 
and  distils  at  a white  heat  (about  1700°)  ; it  burns  with  an  intense  white 
light  when  heated  in  the  air,  and  forms  tin  dioxide.  It  does  not  oxidize 
in  the  air  at  ordinary  temperatures,  and  withstands  the  action  of  many 
bodies,  hence  is  employed  in  tinning  copper  and  iron  vessels  for  house- 
hold use. 

The  most  interesting  of  the  tin  alloys,  besides  bronze  and  soft  solder, 
is  britannia  metal.  It  contains  9 parts  of  tin  and  i i)art  of  antimony, 
and  frequently,  also,  2-3  per  cent,  of  zinc  and  i per  cent,  of  copper. 

Tin  dissolves  in  hot  hydrochloric  acid,  to  stannous  chloride,  with 
evolution  of  hydrogen  gas : 

Sn  + 2HCI  =r  SnCb  -f  211. 

Concentrated  sulphuric  acid,  when  heated,  dissolves  tin,  with  forma- 
tion of  stannous  sulphate. 

Nitric  acid,  depending  upon  the  temperature  and  the  concentration  of 
the  acid,  forms  soluble  stannous  nitrate  or  solid  stannic  nitrate,  which 
separates  and  is  converted  by  the  increasing  dilution  of  the  acid  or  by 
hot  water  into  a basic  salt  and  stannic  acid.  Different  stannic  acids 
result  in  accordance  with  the  conditions  of  experiment.  Anhydrous 
nitric  acid,  HNO3,  does  not  change  tin.  It  dissolves  when  boiled  with 
potassium  or  sodium  hydroxides,  forming  stannates: 

Sn  + 2KOH  + H^O  = K^SnOg  + 2!!,. 

There  are  two  series  of  compounds  : the  stannous,  and  stannic  or  com- 
pounds of  stannic  acid.  The  first  readily  oxidize  to  stannic  compounds. 


STANNOUS  COMPOUNDS. 

Tin  Dichloride  or  Stannous  chloride^  SnCl.^,  results  when  tin  dissolves 
in  concentrated  hydrochloric  acid.  When  its  solution  is  evaporated  it 
crystallizes  with  two  molecules  of  water  (SnCl^  2H3O),  which  it  loses  at 
100°.  It  is  used  in  dyeing,  as  a mordant,  under  the  name  of  tin  salt. 
The  anhydrous  chloride,  obtained  by  heating  the  metal  in  dry  hydro- 


STANNOUS  COMPOUNDS. 


365 


chloric  acid  gas,  fuses  at  250°  and  distils  without  decomposition  at  606°. 
Its  vapor  density  at  900°  agrees  with  the  formula  SnCl2,  while  at  lower 
temperatures  the  molecules  Sn^Cl^  also  seem  to  exist. 

Stannous  chloride  dissolves  readily  in  water.  Its  solution  is  strongly 
reducing,  and  absorbs  oxygen  from  the  air  with  the  separation  of  basic 
stannous  chloride: 

SSnCh  + O + H2O  = 2Sn<Q^  + SnCh* 

In  the  presence  of  hydrochloric  acid,  only  stannic  chloride  is  pro- 
duced. Stannous  chloride  precipitates  mercurous  chloride  or  metallic 
mercury  from  solutions  of  mercuric  chloride  (p.  325).  It  unites  with 
chlorine  to  form  stannic  chloride,  and  with  many  chlorides  to  yield 
double  salts,  e.  g.  : 

SnCl2. 2KCI  and  SnC^ . 2NH,C1. 

It  is  decomposed  by  concentrated  hydrochloric  acid  with  evolution  of 
hydrochloric  acid  gas. 

Tin  Monoxide  or  Sta^inous  oxide,  SnO,  is  obtained  by  heating  its 
hydroxide,  Sn(OH)2,  in  an  atmosphere  of  carbon  dioxide ; it  is  a blackish- 
brown  powder,  which  burns  when  heated  in  the  air,  and  becomes  stannic 
oxide.  Sodium  carbonate  added  to  a solution  of  stannous  chloride  pre- 
cipitates white 

Stannous  staiino-hydrate,  Sn(OH)2: 

SnCl2  + Na2C03  + H2O  = Sn(OH)2  + 2NaCl  + CO2. 

It  is  insoluble  in  ammonium  hydroxide,  but  is  readily  dissolved  by 
potassium  hydroxide.  Upon  slow  evaporation  of  the  alkaline  solution, 
dark  crystals  of  stannous  oxide,  SnO,  separate;  but,  on  boiling  the  solu- 
tion, the  hydrate  decomposes  into  potassium  stannate,  K2Sn03,  which 
remains  dissolved,  and  metallic  tin  : 

2K2Sn02  + H2O  = K2Sn03  + 2KOH  + Sn. 

The  hydroxide  forms  salts  by  its  solution  in  acids.  Stannous  chloride, 
SnClg,  and  stannous  sulphate,  SnSO^,  are  formed  when  tin  is  warmed  with 
concentrated  hydrochloric  or  sulphuric  acid.  The  sulphate  separates  in 
small,  granular  crystals,  when  its  solution  is  evaporated. 

Tin  Monosulphide,  Stannous  sulphide,  SnS,  is  precipitated  from 
stannous  solutions  by  hydrogen  sulphide,  as  a dark-brown  amorphous 
precipitate.  Obtained  by  fusing  tin  and  sulphur  together,  it  is  a lead- 
gray  crystalline  mass.  It  dissolves  in  concentrated  hydrochloric  acid, 
with  liberation  of  hydrogen  sulphide,  and  forms  stannous  chloride.  It 
is  insoluble  in  alkaline  monosulphides,  but,  if  sulphur  be  added  and  the 
solution  boiled,  it  will  dissolve  as  a sulpho-stannate  (p.  367): 

SnS  -f  S + K2S  = K2SnS3. 


366 


INORGANIC  CHEMISTRY. 


STANNIC  COMPOUNDS. 

Tin  Tetrachloride,  Stannic  chloride^  SnCl^,  is  produced  by  the 
action  of  chlorine  upon  lieated  tin  or  stannous  chloride  (see  Lorenz, 
Z.  f.  anorg.  Ch.  10  (1895)  44)-  is  a colorless  licpiid  {Spiritus  fumans 
Libavii^,  fuming  strongly  in  moist  air,  of  specific  gravity  2.27,  and  boil- 
ing at  114°;  its  va})or  density  corres|)onds  to  the  molecular  formula, 
SnCl^.  It  attracts  moisture  from  the  air  and  is  converted  into  a crystal- 
line mass  (butter  of  tin),  SnCl^ -{- 3^^/^  readily  soluble  in  water. 
Boiling  decomposes  the  solution  into  metastannic  acid  (H^SnO,)  and 
hydrochloric  acid  : 

SnCh -f  3 1 BO  _ 11281103  i 4IICI. 

Stannic  chloride  possesses  a salt-like  nature,  and  combines  with  metallic 
chlorides  to  the  so-called  double  salts,  e.  g.,  Sn(d^.  2KCI  and  SnCI^.- 
2NH4CI ; the  latter  compound  is  known  as  pink  sa/t  iind  is  used  in  calico 
printing.  It  also  yields  crystalline  double  salts  with  chlorides  of  the 
metalloids,  c.  g.,  SnCl^.PClg  and  SnCl^.aSCl^.  It  is  not  attacked  by 
even  hot  concentrated  sulphuric  acid. 

Tin  tetrachloride  combines  with  hydrochloric  acid,  forming  HjSnClg  -p  6II2O,  analo- 
gous to  the  chlorplatinic  acid,  H2PtClg  -[-  61  bO.  It  is  formed  when  hydrochloric  acid  gas 
is  conducted  into  a concentrated  solution  of  tin  tetrachloride  in  water.  In  the  cold  it 
solidifies  to  a leafy  crystalline  mass,  melting  at  4 9°-  These  double  compounds  can  also 
be  regarded  as  salts  of  hydrochlorstannic  acid,  e.  g.,  (Nlb)2SnClg. 

Tin  Bromide,  SnBr^,  forms  a white,  crystalline  mass  that  melts  at  30°  and  boils  at 
200°.  It  unites  with  hydrogen  bromide,  forming  H2SnBrg  4 8H.,0,  crystallizing  in 
yellow  needles  and  plates. 

Tin  Iodide,  Snl^,  formed  upon  heating  tin  and  iodine  at  50°,  consists  of  orange-red 
octahedra,  fusing  at  146°  and  boiling  at  295°.  A solution  of  tin  iodide  in  arsenic 
tribromide  (p.  146)  has  the  specific  gravity  3.73  at  15°  ; it  is  therefore  the  heaviest  solu- 
tion specifically  of  which  we  have  knowledge. 

Tin  Fluoride,  SnFp,  is  only  known  in  combination  with  metallic  fluorides  {e.  g., 
K2SnF]g),  which  are  very  similar  to  and  generally  isomorphous  with  the  salts  of  hydro- 
fluosilicic  acid  (K28iFl6)- 

Tin  Dioxide,  Stannic  oxide,  Sn02,  is  found  in  nature  as  tin-stone  or 
cassiterite,  in  quadratic  crystals  or  compact,  brown  masses,  of  specific 
gravity  6.8.  It  is  prepared,  artificially,  by  heating  tin  in  the  air,  and  it 
then  forms  a white  amorjihous  powder,  or  long,  rhombic  leaflets.  It 
may  be  obtained  crystallized,  by  conducting  vapors  of  the  tetrachloride 
and  water  through  a tube  heated  to  redness.  The  dioxide  fuses  in  the 
electric  furnace,  and  is  not  soluble  in  acids  or  alkalies.  When  fused  with 
sodium  and  potassium  hydroxide  it  yields  stannates  soluble  in  water. 

The  dioxide  forms  two  hydrates — the  stannic  acids,  H.^SnOg  and 
H^Sn(b — corresponding  to  the  hydrates  H2CO3,  H2Si03  and  H^CO^  and 
114^104(1).  226).  Each  of  these  two  acids  is  known  in  two  varieties 
which  differ  very  much  from  one  another  in  their  deportment.  This 
double  appearance  of  similarly  constituted  stannic  acids  was  the  first 
instance  of  isomerism  (|).  87).  Berzelius  (1811)  observed  it,  but  he  did 
not  comprehend  its  full  imjiort  until  after  (1817)  the  investigations  of 
Davy  and  Oay  Lussac  had  been  made.  Tin  and  nitric  acid  of  specific 


LEAD. 


367 


gravity  1.35  yield  stannic  nitrate  and  the  /5-  or  vietastannic  acid,  of  the 
formula  H2Sn03  or  H^SnO^,  depending  on  its  method  of  preparation. 
It  is  white,  voluminous  and  insoluble  in  acids.  It  only  dissolves  when 
freshly  made  in  sodium  hydroxide.  An  excess  of  soda  precipitates 
sodium  metastannate  from  this  solution  ; it  dissolves  in  pure  water. 

The  acid  upon  which  hydrochloric  acid  has  been  poured  dissolves  in  pure 
water.  Sulphuric  acid  reprecipitates  it.  a-Stannic  acid — also  H2Sn03  or 
H^SnO^ — is  precipitated  as  a voluminous  white  powder  on  the  addition  of 
alkali  to  solutions  of  the  tetrachloride  or  bromide  which  have  not  stood 
too  long.  It  dissolves  readily  while  moist  in  mineral  acids  and  in  sodium 
hydroxide,  even  when  the  latter  is  in  excess.  Both  acids  separate  on 
heating  their  dilute  hydrochloric  acid  solutions  or  upon  the  addition  of 
ammonia  water.  The  presence  of  tartaric  acid  prevents  the  precipitation 
of  the  a-acid  by  ammonia.  The  a-acid  in  hydrochloric  acid  solution 
gradually  passes  into  the  /5-acid  (Lorenz,  Z.  f.  anorg.  Ch.  9 (1895),  368). 
The  reason  for  this  first  observed  case  of  isomerism  has  not  been  explained. 

Most  of  the  salts  of  stannic  oxide  with  acids,  e.  g , the  sulphate,  are  not 
very  stable,  and  washing  with  warm  water  decomposes  them.  The  phos- 
phate and  arsenate  are  exceptions.  They  are  insoluble  in  water  and  in 
nitric  acid,  which  is  of  importance  in  analysis.  The  metallic  salts  of  the 
stannic  acids  are  more  stable.  The  most  important  of  these  is 
stannate,  Na2Sn03  -(-  3H2O,  which  is  employed  in  calico  printing  under 
the  name  of  preparing  salts.  It  is  produced  upon  a large  scale  by  fusing 
tin-stone  with  sodium  hydroxide.  On  evaporating  the  solution,  it  crystal- 
lizes in  large,  transparent,  hexagonal  crystals. 

Tin  Disulphide,  Stannic  sulphide,  81182,  is  precipitated  as  an  amor- 
phous yellow  powder  by  hydrogen  sulphide  from  acid  solutions  of  stannic 
salts.  If  a mixture  of  tin  filings,  sulphur,  and  ammonium  chloride  be 
heated  it  is  obtained  in  the  form  of  a brilliant  crystalline  mass,  consist- 
ing of  gold-yellow  scales.  It  is  then  called  mosaic  gold,  and  is  applied  in 
bronzing.  Concentrated  hydrochloric  acid  dissolves  the  precipitated  di- 
sulphide, forming  stannic  chloride  ; nitric  acid  converts  it  into  metastan- 
nic  acid.  The  sulphides  and  hydrosulphides  of  the  alkalies  dissolve  tin 
disulphide,  forming  sulphostannates  (see  p.  223).  Sodhun  sulphostannate, 
Na28n83  -(-  2H2O,  crystallizes  in  colorless  octahedra.  Acids  decompose 
the  sulphostannates  with  the  separation  of  tin  disulphide  and  the  evolu- 
tion of  hydrogen  sulphide.  Freshly  precipitated  tin  disulphide  dissolves 
rather  readily  in  aqueous  ammonia  forming  a red  liquid  which  becomes 
colorless  after  standing  in  the  air  for  some  time.  From  this  solution  acids 
precipitate  a white  tin  oxysulphide,  8n82  + 8n80,  soluble  in  aqueous  am- 
monium carbonate. 


3.  LEAD. 

Pb  = 206.9 

I.ead  {Plumbuni)  is  found  in  nature  principally  as  galenite,  Pb8.  The 
other  more  widely  distributed  lead  ores  are  cerussite,  PbCOg,  crocoisite 
(PbCrOJ,  wulfenite  (PbMoOJ  ; also  pyromorphite , PbClz  -|-  3Pb3(POj2- 


368 


INORCJANIC  CHKMISrJ<Y. 


anglesilc,  PbSO^,  and  bournonitr,  PhS.Sb^S.^.  The  countries  of 
importance  in  tlie  industry  of  lead  are  the  Unite(i  States,  Si)ain,  (Ger- 
many, Clreat  Pritain,  and  New  South  Wales,  (lalenite  isthecln’ef  source 
of  lead;  the  i)rocess  of  its  sei)aration  is  very  simple,  d'he  galenite  is 
first  roasted  in  the  air  and  then  strongly  ignited  away  from  it.  In  the 
roasting,  a portion  of  the  lead  sulphide  is  oxidized  to  oxide  and  sulphate  : 


and 


1‘1)S  I 30  = PbO  + SOj 
ri)S  I o^  ^pbscv 


Upon  ignition,  these  two  substances  according  to  Percy  react  with  the 
lead  sulphide  according  to  the  following  expiations: 


and 


2 PbO  + PbS  = 3Pb  + SO., 
PbSO,  -I-  PbS  = 2Pb  + 2SO,.- 


Galenite  may  also  be  fused  with  iron  : PbS  -j-  Fe  = FeS  -|-  Pb.  Iron 
sulphide,  containing  lead  sulphide,  is  used  in  making  sul[)hur  dioxide  for 
the  lead-chamber  process,  d'he  resulting  work- lead  is  contaminated  with 
nearly  all  the  metals  which  occur  in  ores.  They  are  eliminated  by  rather 
complex  methods. 

Pure  metallic  lead  has  a bluish-white  color,  is  very  soft,  and  tolerably 
ductile.  A freshly  cut  surface  has  a bright  luster,  but  on  exposure  to  air 
becomes  dull  by  oxidation.  Its  specific  gravity  is  11.37.  It  melts  at 
about  300°,  and  distils  at  a white  heat  (about  1700°).  It  burns  to  lead 
oxide  when  heated  in  the  air. 

In  contact  with  air  and  pure  water  lead  oxidizes  to  lead  hydroxide, 
Pb(OH)2,  which  is  somewhat  soluble  in  water.  If,  however,  the  water 
contain  carbonic  acid  and  mineral  salts — even  in  slight  quantity,  as  in 
natural  waters — no  lead  goes  into  solution,  but  it  is  covered  with  an 
insoluble  layer  of  lead  carbonate  and  sulphate.  When  much  carbon 
dioxide  is  present  the  carbonate  is  somewhat  soluble  in  water.  This 
behavior  is  very  important  for  practical  purposes,  as  lead  pipes  are  fre- 
quently employed  in  conducting  water.  Lead  also  dissolves  in  ammonia 
and  liine-water  in  the  presence  of  air. 

Sulphuric  and  hydrochloric  acids  have  little  effect  on  the  compact 
metal,  owing  to  the  insolubility  of  its  sulphate  and  chloride;  yet,  if  the 
lead  be  in  the  form  of  powder,  both  acids  will  dissolve  it.  It  forms  lead 
nitrate  with  dilute  nitric  acid.  Zinc,  tin,  and  iron  precipitate  it,  as 
metal,  from  its  solution  ; a strip  of  zinc  immersed  in  a dilute  solution  of 
lead  acetate  is  covered  with  an  arborescent  mass,  consisting  of  shining 
crystalline  leaflets  (lead  tree). 

Alloys. — An  alloy  of  e(]ual  i)arts  of  lead  and  of  tin  fuses  at  186°,  and  is 
used  for  soldering  (soft  solder).  An  alloy  of  4-5  parts  of  lead  and  i ])art 
of  antimony  is  very  hard,  and  answers  for  the  manufacture  of  type  (hard 
lead  or  type-metal).  Shot  are  lead  containing  0.2-0.35  per  cent,  of 
arsenic. 

d'he  usual  lead  compounds  are  constituted  according  to  the  type  PbX.,, 


LEAD. 


369 


and  are  called  plumbic  (p.  357).  Many  of  the  lead  salts  are  isomorphous 
with  those  of  barium  ; the  sulphates  of  both  metals  are  insoluble  in  water. 
In  addition  to  the  salts  of  the  type  PbX2  lead  also  forms  that  of  the  type 
PbX^,  which  are,  however,  very  unstable. 

Lead  Oxide,  PbO,  is  produced  when  lead  is  heated  in  air.  It  melts 
at  a red  heat  and  after  fusion  it  solidifies  to  a reddish-yellow  mass  of 
rhombic  scabs  (litharge).  When  lead  is  carefully  roasted,  or  the  hydrox- 
ide or  nitrate  ignited,  we  obtain  a yellow  amorphous  powder  called  mas- 
sicot. Lead  oxide  hasstrong  basic  properties;  it  absorbs  carbon  dioxide 
from  the  air,  and  imparts  an  alkaline  reaction  to  water  as  it  dissolves  as 
hydroxide.  Like  other  strong  bases  it  saponifies  fats  (lead  plaster).  It 
dissolves  in  hot  potassium  hydroxide,  and  on  cooling  crystallizes  from 
solution  in  yellow  rhombic  or  red  tetragonal  forms,  depending  upon 
existing  conditions  [Geuther,  Ann.  Chem.  219  (1883),  56]. 

Lead  Hydroxide,  Pb(OH)2.  Alkalies  throw  it  out  of  lead  solutions 
as  a white,  voluminous  precipitate. 

It  dissolves  slightly  in  water,  and  absorbs  carbon  dioxide  with  forma- 
tion of  lead  carbonate.  When  heated  to  130°  it  decomposes  into  lead 
oxide  and  water.  It  is  solul)le  in  caustic  potash  or  soda. 

If  lead  or  the  amorphous  oxide  be  heated  to  300-400°,  for  some  time, 
in  the  air,  it  will  absorb  oxygen  and  be  converted  into  a bright  red  pow- 
der, called  red  lead,  or  minium.  Its  compositi(m  corresponds  to  the 
formula  Pb^ ; it  is  considered  a compound  of  lead  monoxide  with  lead 
peroxide : 

Pb,0,  ^ 2PbO  + Pb02. 


When  minium  is  treated  with  dilute  nitric  acid,  lead  nitrate  passes  into 
solution,  while  a dark-brown  amorphous  powder,  lead  peroxide,  Pb02, 
remains. 

This  oxide  is  more  conveniently  obtained  by  the  interaction  of  lead 
carbonate  and  sodium  hypochlorite,  e , if  a solution  of  lead  acetate  be 
mixed  with  an  excess  of  sodium  carbonate  and  chlorine  be  conducted 
into  the  pasty  mass  : 

PbCOg  + Na2C03  + CI2  = Pb02  + 2NaCl  4 2CO2. 

Lead  peroxide  dissolves  in  ice-cold  concentrated  hydrochloric  acid  to 
a reddish-yellow  liquid  containing  lead  teirachloiide,  PbCl^. 

Oxygen  is  disengaged  when  sulphuric  acid  acts  upon  it,  and  lead  sul- 
phate (PbSOJ  is  formed.  When  dry  sulphur  dioxide  is  conducted  over 
it,  glowing  sets  in  and  lead  sulphate  results  : 

Pb02  + SO2  = PbSO,. 

When  ignited  lead  dioxide  breaks  down  into  lead  monoxide  and 
oxygen. 

Lead  dioxide,  when  warmed  with  potassium  hydroxide,  dissolves,  and  on  cooling, 
large  crystals  of  potassinvi  pluvibate,  K2Pb03  -j  3 HgO,  separate  out  ; these  are  perfectly 
analogous  to  potassium  stannate,  K2Sn<')3  4-  3H2O.  An  alkaline  lead  oxide  solution 
added  to  a solution  of  potassium  plumbate  produces  a yellow  precipitate  (Pl^O^  I^O), 


370 


INORGANIC  CHEMISTRY. 


which  loses  water  upon  gentle  wanning,  amt  is  converted  into  red  lead.  Therefore,  the 
latter  must  be  considered  as  the  lead  salt  of  a nonnal  plumbic  acid,  Pb(OIl)^,  which  cor- 
responds to  stannic,  Sn(Oll)4,  and  silicic,  Si(Oll)^,  acitls  : 

Pl).,()^  - rb,rb(),. 

The  calcium  salt,  — calcium  orthoplumhate,  Ca.^I’bC)^, — is  produced  upon  igniting  lead 
peroxide  with  lime  or  calcium  carl)onate,  or  when  lead  oxide  and  lime  are  heated  with 
air  access,  /.  <?.,  the  absorjjtion  of  oxygen.  It  is  a yellowish-red  substance,  which  breaks 
down  when  heated  in  carb(m  dioxide  into  lead  oxide,  calcium  carbonate  and  oxygen  : 

Ca,Pb(),  t 2CO,  = PbO  4-  2CaC03  + (),. 

The  resulting  mixture  of  lead  oxide  and  chalk  regenerates  calcium  plumbate  when  it  is 
heated  in  an  air  current  (Kassner’s  oxygen  methcKl). 

Lead  peroxide  and  other  oxides  of  lead  j)lay  an  important  part  in  charging  accumulators 
with  electricity.  In  the  process  of  charging  the  lead  oxide  at  the  negative  plate  is  decom- 
])osed  into  lead  and  oxygen,  while  at  the  positive  pole  it  absorbs  oxygen  and  becomes 
peroxide.  In  discharging  these  changes  are  reversed. 

Another  oxide,  Pb./).^,  which  is  precipitated  as  a reddish  yellow  powder  on  the  addi- 
tion of  sodium  hydrocldorite  to  an  alkaline  lead  .solution,  is  very  probably  the  lead  salt  of 
metaplumbic  acid,  ll2pb()., : 

Pb203  PbPbO.,. 

Nitric  acid  decomposes  it  into  lead  nitrate  and  peroxide.  It  dissolves  in  cold  hydro- 
chloric acid  without  liberation  of  chlorine  ; this  gas  escapes,  however,  when  the  solution 
is  heated. 

Lead  Tetrachloride,  PbCl^,  was  prepared  by  H.  Friedrich  in  1893 
[Monats.  f.  Ch.  14,  505].  It  is  formed  on  adding  lead  peroxide  to  cold 
concentrated  hydrochloric  acid,  or  upon  conducting  chlorine  into  con- 
centrated hydrochloric  acid  containing  lead  chloride  in  solution  or  in 
suspension.  On  adding  ammonium  chloride  to  the  clear  liquid  the 
double  salt,  PbCl^.  2NHCI,  separates  in  yellow  isometric  crystals.  It  is 
decomposed,  when  added  to  ice-cold  concentrated  sulphuric  acid,  into 
ammonium  sulphate,  hydrochloric  acid  and  lead  tetrachloride,  which 
collects  under  the  sulphuric  acid  in  the  form  of  a yellow  oil. 

Lead  tetrachloride  becomes  crystalline  at — 15°.  It  decomposes  gradu- 
ally at  the  ordinary  temperature  into  lead  chloride  and  chlorine.  The 
decomposition  is  explosive  at  105°;  its  specific  gravity  at  0°  is  3.18.  It 
fumes  in  moist  air.  It  is  soluble  in  water.  Concentrated  solutions  when 
warmed  liberate  chlorine  and  lead  chloride  separates  ; in  dilute  solutions 
lead  peroxide  is  precipitated.  Sodium  and  potassium  hydroxides  act 
similarly.  The  group  similarity  with  carbon,  silicon,  germanium  and 
tin  exi)resses  itself  in  this  chloride  ; the  chlorides  of  these  other  members 
are  liquids  at  the  ordinary  temperature. 

Lead  Chloride,  PbCl^,  separates  as  a white  precipitate,  when  hydro- 
chloric acid  is  added  to  the  solution  of  a lead  salt.  It  is  almost  insoluble 
in  cold  water;  from  hot  water,  of  which  it  requires  30  parts  for  solution, 
it  crystallizx's  in  white,  shining  needles.  It  melts  at  about  500°  and  solidi- 
fies to  a horn-like  mass.  It  is  volatile  at  a white  heat;  its  vapor  density 
corresponds  to  the  formula  P1)C1.2. 

Lead  Iodide,  Pbl2,  is  thrown  down  as  a yellow  precipitate  from  lead 
solutions  by  j)otassium  iodide  ; it  crystallizes  from  a hot  solution  in  shin- 
ing, yellow  leaflets,  melting  at  383°.  It  is  difficultly  soluble  in  water. 


LEAD.  371 

Lead  tetraiodide,  Pbl4,  and  tetrabromide,  PbRr,,  are  known  in  the  form  of  double 
salts  with  salts  of  organic  bases  (Classen  and  Zahorski). 


Lead  Nitrate,  Pb(N03).^,  obtained  by  the  solution  of  lead  in  dilute 
nitric  acid,  crystallizes  in  regular  octahedra  (isomorphous  with  barium 
nitrate)  and  dissolves  in  two  jtarts  of  water  at  the  ordinary  temperature. 
It  melts  at  a red  heat,  and  is  decomposed  into  lead  monoxide,  nitrogen 
dioxide,  and  oxygen.  When  boiled  with  lead  oxide  and  water,  it  is  con- 
verted into  the  basic  nitrate,  which  separates  in  white  needles. 

Lead  Sulphate,  PbSO^,  occurs  in  nature  as  ang/esiie,  in  rhombic 
crystals,  isomorphous  with  barium  sulphate.  It  is  precipitated  from  lead 
solutions  as  a white  crystalline  mass  by  sulphuric  acid.  It  dissolves  with 
difficulty  in  water,  more  readily  in  concentrated  sulphuric  acid,  very  easily 
in  concentrated  sodium  hydroxide  (distinction  from  barium  sulphate). 
When  ignited  with  carbon,  it  is  decomposed  according  to  the  following 
equation  : 

PbSO^  + 2C  = PbS  + 2CO.,. 

Lead  Carbonate,  PbCOg,  occurs  as  cerussite  in  nature.  It  is  pre- 
cipitated by  ammonium  carbonate  from  lead  nitrate  solutions.  Potassium 
and  sodium  carbonates  precipitate  basic  carbonates,  the  composition  of 
which  varies  with  the  temperature  and  concentration  of  the  solution. 
A similar  basic  salt,  whose  composition  agrees  best  with  the  formula : 

2PbC03.  Pb(OH)2* 

is  prepared  on  a large  scale  by  the  action  of  carbon  dioxide  upon  lead 
acetate.  It  bears  the  name  white  lead. 

White  lead  was  formerly  manufactured  by  what  is  known  as  the  Dutch  process.  Leaden 
plates  rolled  up  into  a spiral  were  moistened  in  earthenware  pots  with  acetic  acid,  and  then 
covered  with  manure  and  permitted  to  stand  undisturbed  for  some  time.  In  this  way  the 
action  of  the  acetic  acid  and  air  upon  the  lead  produced  a basic  acetate,  which  the  carbon 
dioxide,  evolved  from  the  decaying  manure,  converted  into  basic  lead  carbonate.  At 
present  it  is  prepared  by  dissolving  litharge  in  acetic  acid,  and  converting  the  resulting 
basic  acetate  into  a carbonate  by  conducting  carbon  dioxide  into  it. 

White  lead  is  employed  for  the  manufacture  of  white  oil  colors.  As  it  is  poisonous, 
and  blackened  by  the  hydrogen  sulphide  of  the  air' (formation  of  lead  sulphide),  it  is 
being  replaced  more  and  more  by  zinc  white  and  permanent  white  (BaSO^). 

Lead  Sulphide,  PbS,  occurs  crystallized  in  metallic,  shining  cubes 
and  octahedra.  Hydrogen  sulphide  precipitates  it  as  an  amorphous  black 
y^owder.  It  is  insoluble  in  dilute  acids.  Hydrogen  sulphide  preciydtates 
red  chloro-sulyffiide,  PbS  . PbClg,  from  solutions  of  lead  salts  containing 
much  strong  hydrochloric  acid. 

The  soluble  lead  compounds  are  very  poisonous.  They  have  a sweet- 
ish, astringent  taste.  They  are  readily  recognized  by  the  following  reac- 


*Its  chemical  structure  may  be  expressed  by  the  following  formula  : 

HO-Pb-COg-Pb-COg-Pb-OH. 


372 


INORGANIC  CHEMISTRY. 


tions  : sulpliuric  acid  i)recipitates  wliite  lead  siili)hale,  wliich  is  soluble 
in  sodium  hydroxide;  from  this  solution  hydrogen  sulphide  ])reci])itates 
black  lead  sulphide.  Potassium  iodide  [)recii)itates  yellow  lead  iodide. 


4.  BISMUTH. 

\V\  = 208.5. 

Bismuth  constitutes  a natural  group  with  antimony,  arsenic,  phosphorus 
and  nitrogen.  We  observed  that,  with  increase  of  atomic  weight,  the 
metalloidal  character  of  the  lower  members  becomes  more  metallic  ; the 
acid  nature  of  the  oxides  becomes  basic.  Antimony  oxide  (SI)./),)  is  a 
base,  while  the  higher  oxide,  Sb.p^^,  represents  an  acid  anhydride.  In 
bismuth,  the  metallic  nature  attains  its  full  value.  This  is  manifest  in 
its  inability  to  unite  to  a volatile  compound  with  hydrogen  and  in  the 
basic  oxyhydrate  BiO(OH).  Bismuth  trioxide  is  a base,  and  the  pent- 
oxide  j)ossesses  a very  feeble  acid  character ; the  latter  behaves  more  like 
a metallic  peroxide. 


Bismuth  usually  occurs  native,  and  in  combination  with  sulphur,  as 
bismuthinite  ; vvitli  tellurium  as  tetradymite ; also  as  bisjnuih  ocher ^ Bia^g? 
and  in  many  cobalt,  nickel  and  silver  ores. 

Bismuth  is  obtained  principally  from  Saxony,  and  in  London  from  Bolivian  and  Aus- 
tralian ores.  In  Saxony  the  roasted  ores  are  extracted  with  concentrated  hydrochloric 
acid,  and  the  solution  then  diluted  with  water,  when  the  oxychloride  is  precipitated.  It 
is  redi.ssolved  in  hydrochloric  acid  and  freed  from  iron  by  precipitation  with  water,  and 
then  melted,  together  with  coal,  lime,  and  slag,  in  graphite  crucibles  to  obtain  the 
bismuth. 

To  obtain  the  metal,  the  sulphide  is  roasted  in  the  air,  and  the  result- 
ing oxide  reduced  with  charcoal. 

Bismuth  is  a reddish-white  metal,  of  specific  gravity  9.9.  It  is  brittle 
and  may  be  easily  pulverized.  It  crystallizes  in  rhombohedra.  It  fuses 
at  265°  and  distils  at  a white  heat  (about  1300°).  It  does  not  change  in 
the  air  at  ordinary  temperatures.  When  heated  it  burns  to  bismuth 
oxide,  Bi203.  It  is  insoluble  in  hydrochloric  acid,  but  dissolves  in  boil- 
ing sulphuric  acid  with  formation  of  sulphate  of  bismuth,  and  the  evolu- 
tion of  sulphur  dioxide.  Nitric  acid  dissolves  it  readily  in  the  cold. 

Water  decomposes  bismuth  solutions  in  the  same  manner  as  those  of 
antimony;  insoluble  basic  salts  are  precipitated,  but  these  are  not  dis- 
solved by  tartaric  acid  (distinction  from  antimony). 

Bismuth  Chloride,  BiCl^,  arises  from  the  action  of  chlorine  upon 
heated  bismuth,  and  by  the  solution  of  the  metal  in  aqua  regia.  It  is  a 
white  crystalline  mass,  which  fuses  at  220°  and  boils  at  about  447°.  It 
deliquesces  in  the  air.  Water  renders  its  hydrochloric  acid  solution  turbid, 
a white,  crystalline  ])recipitate  of  bismuth  oxychloride,  BiOCl,  separating 
at  the  same  time  ; 

BiCIg  + lIjO  = BiOCl  T 2HCI. 

The  metalloidal  character  of  bismuth  is  indicated  bv  this  reaction. 


CHROMIUM  GROUP. 


373 


The  compounds  BiBr^  (orange-yellow)  and  Bilg  (black  or  dark  brown) 
are  very  similar  to  bismuth  chloride.  All  three  combine  with  many 
metallic  haloid  salts  to  form  double  halogen  derivatives. 

Halogen  derivatives  of  quinquivalent  bismuth  are  unknown. 

Bismuth  Oxide,  Bi203,  prepared  by  burning  bismuth  or  heating  the 
nitrate,  is  a yellow  powder,  insoluble  in  water  and  the  alkalies. 

Normal  bismuth  hydroxide,  Bi(OH)3,  is  not  known  in  a free  state. 
Potassium  hydroxide  added  to  a bismuth  solution  precipitates  a white 
amorphous  7)ietahydrate,  BiO.  OH. 

Chlorine  conducted  through  a concentrated  potassium  hydroxide  solu- 
tion in  which  bismuth  oxide  is  suspended  precipitates  red  bismuthic 
acid  (HBiOg  or  H.2Bi20g),  which,  when  gently  heated,  becomes  Bi203, 
bis7iiuthic  oxide.  Strong  ignition  converts  the  latter  into  bismuth  tri- 
oxide and  oxygen;  hydrochloric  acid  dissolves  it  to  bismuth  chloride, 
with  liberation  of  chlorine. 

Bismuth  Nitrate,  Bi(N03)3,  is  obtained  by  the  solution  of  bismuth 
in  nitric  acid,  and  crystallizes  with  five  molecules  of  water  in  large,  trans- 
parent tables.  In  a little  water  it  dissolves  without  any  change;  much 
water  renders  it  turbid,  owing  to  the  precipitation  of  white,  curdy  basic 
CNO,  CNO, 

salts:  Bi  < NO^and  Bi  < OH.  The  precipitate  is  emploved  in  medicine 

(oh  (oh 

under  the  name  of  Bis7}iuthu7n  sub7titricu77i  {sub7iitrafe').  It  frequently  con  - 
tains  tellurium. 

Bismuth  Sulphate,  Bi2(S04)3,  is  formed  when  bismuth  dissolves  in 
sulphuric  acid.  It  crystallizes  in  delicate  needles.  Bismuth  Sulphide, 
1^283,  occurring  as  bismuthinite,  is  thrown  down  as  a black  i)recipitate 
from  bismuth  solutions  by  hydrogen  sulphide.  Unlike  antimony  and 
arsenic  sulphides,  it  does  not  form  sulpho-salts. 

The  alloys  of  bismuth  are  nearly  all  readily  fusible.  An  alloy  of  4 
parts  of  bismuth,  i part  of  cadmium,  i part  of  tin,  and  2 ])arts  of  lead,  fuses 
at  65°  (Wood’s  metal).  The  alloy  of  9 parts  of  bismuth,  i part  of  lead 
and  I part  of  tin  (Rose’s  metal)  fuses  at  94°. 

Alkaline  stannous  solutions  precipitate  metallic  bismuth  from  solutions 
of  its  salts : 

2BiO . OH  + 3X2800,  = 2Bi  + 3X28003  + H2O. 


CHROMIUM  GROUP. 

We  observed  that  a group  of  more  metallic  analogous  elements  : titan- 
ium, zirconium  and  thorium  attached  itself  to  the  metalloidal  elements, 
carbon,  silicon  and  tin  (p.  238);  and  further  that  their  was  an  analogous 
group  of  more  metallic  elements : vanadium,  niobium  and  tantalum,  cor- 
responding to  the  metalloidal  group  of  phosphorus  (|).  226).  We  now 
meet  a group  of  metals,  consisting  of  chromium,  molybdenum,  tungsten, 
and  probably  uranium,  that  bears  a like  relation  to  the  elements  of  the 


374 


INORGANIC  CHEMISTRY. 


sulphur  group  (see  Periodic  System  of  the  Elements).  The  resemblance 
of  these  elements  to  sulphur  and  its  analogues  is  plainly  manifest  in  their 
highest  oxygen  compounds  X()^,  and  their  derivatives  (see  also  Man- 
ganese). As  the  elements  of  the  sulphur  group  in  their  highest  oxygen 
compounds  are  sexivalent,  so  chromium,  molybdenum,  tungsten  and 
uranium  form  acid  oxides:  CrO,,  M0O3,  UO3.  Many  of  the  salts 

corresponding  to  these  are  very  similar  to  and  isomorphous  with  the  salts 
of  suli)huric  acid.  Sodium  chromate,  like  sodium  sulphate,  crystallizes 
with  ten  molecules  of  water  ; the  potassium  salts  of  both  groups  form  iso- 
morphous mixtures;  their  magnesium  salts,  as  well  as  that  of  tungstic 
acid,  have  the  same  constitution  : 

MgSO,  + 7H2O  and  MgCrO^  + 7II2O. 

Corresponding  to  the  acid  oxides  are  the  chlorine  derivatives: 

SO2CI2,  CrO,Cl2,  M0O2CI2,  MoOCl,,  WOCl,  and  WClg, 

which  are  perfectly  analogous,  so  far  as  chemical  deportment  is  con- 
cerned. 

The  most  important  basic  oxide  of  chromium  is  its  sesquioxide.  This 
affords  salts  with  the  acids,  and  they  are  perfectly  similar  to  those  of  the 
sesquioxides  of  iron  (Fe203),  manganese  (Mn203),  and  aluminium  (AI2O3) 
(P-  33°)- 

II 

Finally,  compounds  of  chromium,  CrX2,  are  known  in  which  the  metal 
figures  as  a dyad.  These  so-called  chromoiis  compounds  are  very  unstable, 
and  are  oxidized  by  the  air  into  chromic  compounds.  In  this  respect 
they  also  resemble  the  ferrous  derivatives  FeX2  (p.  398). 

Chromium  compounds  containing  more  oxygen  than  the  chromates  are 
found  in  the  salts  of  perchromic  acid ; however,  nothing  definite  is  known 
of  their  constitution  at  present.  The  same  is  true  of  tungsten  and  molyb- 
denum (pp.  381,  383,  384). 

Salts  of  molybdenum  and  tungsten,  corresponding  to  the  states  of 
lowest  oxidation,  are  not  known,  because  these  metals  occur  as  hexads 
in  most  of  their  derivatives.  Uranium,  which  has  the  highest  atomic 
weight  of  the  group,  shows  some  variations  from  its  analogues. 


1.  CHROMIUM. 

Cr  = 52.1. 

Chromium  is  found  principally  as  chromite  in  nature.  This  is  a com- 
bination of  chromic  oxide  with  ferrous  oxide,  Cr203.Fe0,  and  occurs  in 
North  America,  Sweden,  Hungary,  and  in  large  quantities  in  the  Urals. 
Crocoisite,  or  lead  chromate  (PbCrOJ,  is  not  met  with  so  frequently, 
(diromite  is  used  almost  exclusively  for  the  preparation  of  all  other 
chromium  derivatives,  as  it  is  first  converted  into  potassium  chromate 
by  fusion  with  potassium  carbonate  and  nitrate  (p.  379) 

Metallic  chromium  was  first  obtained  free  from  cai  bon  and  fused  in  large 


CHROMIC  COMPOUNDS.  375 

quantities  in  1898  by  Goldschmidt,  who  ignited  a mixture  of  chromic 
oxide  and  aluminium  powder  (p.  347)  : 

Itis  bright  gray  in  color,  very  brilliant,  and  on  cleavage  faces  shows  strong 
crystallization.  It  is  very  hard,  extremely  difficult  to  melt  and  very  stable 
in  the  air.  W.  Hittorf  claims  that  chromium  behaves  like  a noble  metal  at 
the  ordinary  temperature  in  that  it  is  inactive  electromotively  and  in  the 
tension  series  arranges  itself  with  platinum.  Hydrochloric  acid  at  0° 
has  no  action  upon  it  and  when  it  is  made  the  anode  it  dissolves  in  the 
acid  as  chromic  acid.  At  more  elevated  temperatures  it  decomposes 
hydrochloric  acid  with  the  liberation  of  hydrogen  ; and  in  an  electro- 
motive respect  follows  zinc.  It  dissolves  as  chromous  chloride. 

Four  series  of  chromium  compounds  are  known  : chromous  (CrX2), 
chromic  (CrX.^),  the  derivatives  of  chromic  acid,  called  chromates,  and 
the perchromates.  All  chromium  compounds  are  brightly  colored,  hence 
the  name  chromium  (from  color). 


CHROMOUS  COMPOUNDS. 

These  are  very  unstable,  and  by  oxidation  pass  readily  into  chromzV  compounds.  Like 
ferrous  salts,  they  are  produced  by  the  reduction  of  the  higher  oxides.  The  following  may 
be  mentioned : Chromous  Chloride,  CrClg.  This  is  obtained  by  heating  chromic 
chloride,  CrCl.^,  in  a stream  of  hydrogen  or  by  dissolving  chromium  in  hydrochloric  acid. 
It  is  a white  crystalline  powder.  It  volatilizes  without  decomposition.  At  1300-1600° 
the  vapor  density  corresponds  to  a mixture  of  the  molecules  CrClg  and  CrgCh.  It  dis- 
solves in  water  with  a blue  color ; the  solution  absorbs  oxygen  with  avidity,  and  becomes 
green  in  color  [Ann.  Chem.  228  (1885),  113].  The  alkalies  precipitate  yellow  chro- 
mous hydroxide,  Cr(OH)2,  from  it.  This  is  readily  oxidized.  When  heated  it  parts 
with  hydrogen  and  water  and  becomes  chromic  oxide  : 

2Cr(OH)2  = Cr203  + H2  + H2O. 


CHROMIC  COMPOUNDS. 

Chromic  Chloride,  CrClg,  like  AICI3,  is  obtained  by  ignition  of  the 
oxide  and  charcoal  or,  better,  the  metal  in  a current  of  chlorine.  When 
raised  to  a red  heat  in  this  condition  it  sublimes  in  shining  violet  leaflets, 
which  are  transformed  into  chromic  oxide  by  ignition  in  the  air.  Its 
vapor  density  at  1200-1300°  corresponds  to  the  formula  CrClg;  it  vapor- 
izes very  slowly  below  1000°. 

Pure  chromic  chloride  only  dissolves  in  water  after  long-continued 
boiling;  if,  however,  it  contains  traces  of  chromous  chloride,  CrCl2,  it 
dissolves  readily  at  ordinary  temperatures.  Green  crystals  of  CrClg  -|- 
6H2O  separate  from  the  green  solution  on  evaporation  ; these  deliquesce 
in  the  air.  The  same  crystals  may  be  obtained  from  solutions  of  chromic 


376 


INORGANIC  CHEMISTRY. 


hydroxide,  Cr(OH)3,  in  hydrocliloric  acid.  When  they  are  dried  inter- 
mediate oxychlorides,  CrCl2(OH)  and  Cr2Cl(OI  I)^,  are  formed  ; finally 
Cr(OH)3  remains. 

Chromic  Hydroxide,  Cr(OH),,.  It  is  i)recii)itated  by  ammonium 
hydroxide  from  chromic  solutions  as  a voluminous  bluish  gray,  hydrous 
mass.  The  green  preciju tales  i)roduced  by  sodium  and  i)Otassium  hydrox- 
ides contain  alkali  that  cannot  be  removed  even  by  boiling  water. 
They  dissolve  readily  with  an  emerald-green  color,  in  an  excess  (T  jiotas- 
sium  or  sodium  hydroxide  (slightly  in  ammonia),  but  they  are  reprecij)!- 
tated  upon  boiling  their  solutions.  When  chromic  hydroxide  is  healed 
to  200°  in  a current  of  hydrogen,  the  i)roduct  is  the  hydroxide,  CrC). Oil, 
which  is  a grayish-blue  i)owder,  insoluble  in  dilute  hydrochloric  acid. 

Chromic  Oxide,  Cr^O^,  is  a green,  amorphous  powder.  It  is  also 
formed  by  the  ignition  of  chromium  trioxide: 

2Cr03  -=  C1-2O3  -f-  3O, 
or  of  ammonium  dichromate: 

(NIb),Cr.A  = Cr203  + 41^0  + N^. 

It  may  be  obtained  in  black,  hexagonal  crystals,  by  conducting  the 
vapors  of  the  oxychloride  (p.  380)  through  a tube  heated  to  redness: 

2Cr02Cl2  = Cr203  + 201^  + O. 

Ignited  chromic  oxide  is  insoluble  in  acids.  When  fused  with  silicates, 
it  colors  them  emerald-green,  and  serves,  therefore,  to  color  glass  and 
porcelain. 

Giiignef  s green  is  a beautifully  green-colored  chromium  hydroxide, 
which  is  applied  as  a paint.  It  is  obtained  by  igniting  a mixture  of  one 
part  of  potassium  dichromate  with  three  parts  of  boric  acid  ; after  treating 
the  mass  with  water,  which  dissolves  potassium  borate,  there  remains  a 
green  powder,  the  composition  of  which  corresponds  to  the  formula : 

Cr20(0H)^. 


The  predominating  properties  of  chromic  oxide  are  basic,  as  it  readily 
affords  salts  with  acids;  yet  its  basic  nature,  like  that  of  all  sequioxides, 
is  but  slight,  so  that  it  does  not  afford  salts  with  weak  acids  (p.  374). 
In  addition  to  all  this  it  possesses  a slightly  acidic  character,  and  metallic 
salts  are  derived  from  it,  generally  from  the  hydroxide,  CrO.OH,  which 
are  analogous  to  the  aluminates  (p.  351).  Such  salts — chromites — like 
(OO . 0)2Mg  and  (CrO  . can  be  obtained  crystallized  in  regular 

octahedra  by  fusing  chromic  oxide  with  metallic  oxides  and  boron  trioxide 
(as  flux).  The  mineral  chromite  is  such  a salt : 

Cr2O3.Fe0-^  (CrO.O)2Fe. 

Chromium  Sulphate,  Cr2(SO,)3,  is  obtained  by  dissolving  the 
hydroxide  in  concentrated  sulphuric  acid.  'I'he  solution,  green  at  first. 


DERIVATIVES  OF  CHROMIC  ACID. 


377 


becomes  violet  on  standing,  and  deposits  a violet-colored  crystalline 
mass.  This  may  be  purified  by  solution  in  water  and  precipitation  by 
alcohol.  This  salt  crystallizes  from  very  dilute  alcohol  in  bluish-violet  octa- 
hedra  containing  fifteen  molecules  of  water.  If  the  aqueous  solution  of 
the  violet  salt  be  heated,  it  assumes  a green  color,  because  the  salt  breaks 
down  into  free  acid  and  a basic  salt  which,  upon  evaporation,  separates 
as  a green  amorphous  mass,  soluble  in  alcohol.  When  the  green  solution 
stands,  it  reverts  to  the  violet  of  the  neutral  salt.  The  other  chromic 
salts,  the  nitrate  and  the  alum,  behave  in  a similar  manner. 

It  is  very  probable  that  in  the  green  solutions  the  chromium  exists  as 
a part  of  peculiar  “complex”  acids:  chrom-sulphuric  acid,  chrom-nitric 
acid ; it  no  longer  appears  as  ion,  but  as  a part  of  the  latter.  Hence  the 
salts  of  green  solutionsdo  not  show  the  reactions  of  salts  of  chromic  oxide 
nor  of  sulphuric  acid. 

Chromium  sulphate  forms  double  salts  with  the  alkaline  sulphates — the 
chromium  alums  (p.  352). 

Potassium  Chromium  Alum,  K2SO^ . Cr2(SOj3 24H2O,  crys- 
tallizes in  large,  dark-violet  octahedra.  It  is  most  conveniently  ])repared 
by  conducting  sulphur  dioxide  through  a solution  of  potassium  dichromate 
containing  sulphuric  acid  ; 

K2Cr20,  + H2SO,  + 3SO2  = Cr2(SOd3 . K2SO,  + H2O. 

At  80°  the  violet  solution  of  the  salt  becomes  green,  and  on  evapora- 
tion yields  an  amorphous  green  mass. 

As  chromium  hydroxide  possesses  only  a slightly  basic  nature,  salts 
with  weak  acids,  like  carbonic,  sulphurous,  and  hydrosulphuric  acids  (see 
Aluminium,  p.  325)  do  not  exist.  Therefore,  the  alkaline  carbonates  and 
sulphides  precipitate  chromium  hydroxide  from  solutions  of  chromium 
salts : 

Cr2(SOd3  + 3Na3C03  -f  3H2O  = 2Cr(OH)  + 3Na3SO,  + 3CO2 

and 

-1-  3(NHd2S  + 6H2O  = 2Cr(OH)3  + 3(NHd2SO,  + 3H2S. 

Ammonium  sulphide  produces  a black  precipitate,  chromous  sulphide, 
CrS,  in  solutions  of  chromous  salts. 


DERIVATIVES  OF  CHROMIC  ACID. 

In  its  highest  oxygen  derivative,  CrOg,  chromium  possesses  a complete 
metalloidal,  acid-forming  character.  Chromic  acid,  HgCrO^,  is  perfectly 
analogous  to  sulphuric  acid,  H2SO^,  but  has  not  been  obtained  free,  since 
when  liberated  from  its  salts  it  at  once  breaks  down  into  the  oxide  and 
water  : 

HgCrO^  = CrOg  + H2O. 

The  chromates  are  often  isomorphous  with  the  corresponding  sulphates 
(P*  379)-  Polychromaies  also  exist,  and  are  derived  from  polychromic 
32 


INORGANIC  CHEMISTRY. 


acids  produced  by  tlie  condensation  of  several  molecules  of  the  inninal 
acid  (see  Disulphuric  Acid,  p.  193)  : 

Il^CrO^,  IbCr./).,  I etc. 

Chromic  acid.  Dichromic  acid.  Trichromic  acid. 


The  constitution  of  their  salts  is  e.xprcssed  by  the  following  formulas; 


Cr(),< 


OMe 

OMe 


Cr(),< 

Cr().,< 


OMe 

O 

( )Me 


Cr(),,< 

CrO.,< 

CrO,< 


OMe 

O 

o 

OMe. 


The  free  polychromic  acids  are  not  known,  because  as  soon  as  they  are 
sejiarated  from  their  salts,  they  immediately  break  down  into  the  acid 
oxide  and  water  : 


= 3CrO,  + II.p. 


The  polychromates  are  fretpiently,  but  incorrectly,  called  salts ; 
true  acid  or  ])rimary  salts,  in  which  only  one  hydrogen  atom  is  replaced 
by  metal  (KHCrOJ,  are  unknown  for  chromic  acid. 

The  salts  of  normal  chromic  acid  are  mostly  yellow  colored,  while  the 
polychromates  are  red.  The  latter  are  produced  from  the  former  by  the 
action  of  acids  ; 


2K,CrO,  + 2IIXO3  = K,Crp^  -f  2KNO3  + 

Conversely,  by  the  action  of  the  alkalies,  the  polychromates  pass  into  the 
normal  salts : 

K2Ci'207  + 2KOII  = 2K2Cr04  -f  IbO. 

The  formation  of  the  polychromates  maybe  explained  as  follows : The 
chromic  acid  liberated  from  its  salts  by  stronger  acids  breaks  down  into 
water  and  the  acid  oxide,  which  combines  with  the  excess  of  the  normal 
chromate  : 

K2Cr04  -h  CrOg  = K2Cr20^. 

When  there  is  an  excess  of  acid  the  anhydride  (CrOg)  is  set  free. 

Chromium  Trioxide,  or  Chromic  acid  a?thydride  {Acidnm  chro- 
f?iicu?n),  CrOg,  consists  of  long,  red,  rhombic  needles  or  prisms,  ob- 
tained by  adding  sulphuric  acid  to  a concentrated  potassium  dichromate 
solution.  The  crystals  are  readily  soluble  in  water  and  deliquesce  in 
the  air  if  not  perfectly  free  from  sulphuric  acid.  When  heated,  they 
blacken,  melt,  and  at  about  250°  decompose  into  chromic  oxide  and 
oxygen  : 

2Cr03  = Ci-gOg  + 3O. 

It  volatilizes  without  decomposition  in  slight  quantities  and  this  even 
occurs  below  its  jjoint  of  fusion. 

Chromium  trioxide  is  a poweiful  oxidizing  agent,  and  destroys  organic 
matter;  hence  its  solution  cannot  be  filtered  through  paiier.  When  alco- 
hol is  iioured  on  the  crystals,  detonation  takes  })lace,  the  alcohol  burns, 
and  green  chromic  oxide  remains.  By  the  action  of  acids,  e.  g.,  sul- 
phuric, the  trioxidc  deports  itself  like  a peroxicle;  oxygen  escapes  and  a 


DERIVATIVES  OF  CHROMIC  ACID. 


379 


chromic  salt  results.  When  heated  with  concentrated  hydrochloric  acid 
chlorine  is  evolved  : 

2Cr03  + 12HCI  = Ci-jClg  + 6II2O  + 3CI2. 

Reducing  substances,  like  sulphurous  acid  and  hydrogen  sulphide,  con- 
vert chromic  acid  into  oxide  : 

2Cr03  -|-  3112^  — ^^2^3  ~f'  “1“  3S, 

2Cr03  + 3SO2  + 3H2O  ==  Cr203  + 3H2SO^  = €12(80^3  + 3H2O. 

Alcohol  acts  similarly,  being  oxidized  to  aldehyde  : 

^CrOg  + 3H2SO4  -j-  3C2HgO  = Crg (804)3  + 6H2O  3C2H4O. 

Potassium  Chromate,  K2Cr04,  is  obtained  by  adding  potassium 
hydroxide  to  potassium  dichromale.  It  forms  yellow  rhombic  crystals, 
isomorphous  with  potassium  sulphate  (K2SO4) ; isomorphous  mixtures  crys- 
tallize out  from  the  solution  of  the  two  salts. 

Neutral  alkali  chromates  are  always  produced  when  any  chromium  com- 
pound is  fused  with  an  alkaline  carbonate  and  some  oxidizing  agent  (salt- 
peter, potassium  chlorate).  The  addition  of  the  oxidant  is  unnecessary 
if  the  oxygen  of  the  air  can  act  sufficiently  upon  the  fusion.  Chromic 
iron  ore  is  roasted  in  reverberatory  furnaces  with  soda  and  lime.  Neutral 
sodium  chromate,  ferric  oxide,  and  calcium  carbonate  are  the  products : 

2^203.  FeO  -j-  4Na2C03  -f-  4CaO  -j-  7O  = 4Na2Cr04  4CaC03  -f-  ^^203- 

The  neutral  chromate  is  extracted  with  water.  When  this  solution  is 
concentrated  and  mixed  with  sulphuric  acid  anhydrous  sodium  sulphate 
separates  even  from  the  hot  liquid,  from  which,  on  further  evaporation. 
Sodium  Dichromate,  Na2Cr20^  -f  2H2O,  deposits  in  red  crystals  on 
cooling : 

2Na2Cr04  + H28O4  = Na2Cr207  + Na2S04  + H2O. 

All  other  chromium  preparations  are  made  from  this  salt  [see  Hausser- 
mann,  Jahrb.  f.  Ch.  (1891),  1,  327].  Chromates  are  also  produced  if 
chromium  or  its  alloys  are  made  the  anode  in  an  alkaline  bath. 

Potassium  Dichromate,  or  Bichromate  of  potash,  called 

acid  potassium  chromate,  is  manufactured  on  a large  scale  by  trans- 
posing the  sodium  salt  with  potassium  chloride : 

Na2Cr20.j  -f  2KCI  = K2Cr20^  -f  2NaCl. 

The  great  difference  in  solubility  in  water  renders  the  separation  of  the 
sodium  chloride  and  potassium  dichromate  an  easy  matter. 

Potassium  dichromate  (called  red  chromate  of  potash  in  commerce) 
crystallizes  in  large,  red,  triclinic  prisms,  soluble  at  ordinary  tempera- 
tures, in  ten  parts  of  water.  When  heated,  the  salt  fuses  without  change  ; 
at  a very  high  heat  it  decomposes  into  potassium  chromate,  chromic 
oxide,  and  oxygen : 

2K2Cr207  = 2K2CrO^  -j-  Cr203  3O. 


380 


INORGANIC  CHEMISTRY. 


When  the  salt  is  warmed  with  sulphuric  acid,  oxygen  escapes  and  potas- 
sium chromium  alum  is  produced  : 

lY2Cr207  4"  411280^  — Cr2(S()j)3.  1^280^  -j-  41120  -f-  3O. 

This  reaction  answers  for  the  preparation  of  iierfectly  pure  oxygen. 
P\irther,  the  mixture  is  made  use  of  in  laboratories,  as  an  oxidizing  agent. 

Sodium  Chromate,  Na2CrO^  -f-  10H2O,  forms  deliciuescent  crys- 
tals, and  is  analogous  to  ('dauber’s  salt  (Na2S()^  -f-  10H2O).  Barium 
and  Strontium  Chromates,  BaCrC)^  and  SrCrO^,  are  almost  insoluble 
in  water.  Calcium  Chromate,  CaCrO^,  dissolves  with  difficulty  in 
water,  and  crystallizes  like  gypsum  with  two  molecules  of  water,  d'he 
7?iagnesiu?n  salt,  MgCrO^  -|-  7H2^>  dissolves  readily  and  corresponds  to 
Ep.som  salts,  MgSC)^ -j-  7H2().  I'he  chromates  of  the  heavy  metals  are 
insoluble  in  water,  and  are  obtained  by  transposition. 

Lead  Chromate,  PbCrO^,  is  obtained  by  the  precipitation  of 
soluble  lead  salts  with  potassium  chromate.  It  is  a yellow  amorphous 
powder  which  serves  as  a yellow  paint — chrome  yellow.  When  heated  it 
melts  undecomposed,  and  solidifies  to  a brown,  radiating  crystalline 
mass.  It  oxidizes  all  the  carbon  compounds  at  a red  heat,  and  is  there- 
fore used  in  their  analysis.  In  nature  lead  chromate  exists  as  crocoisite. 


Chromic  Acid  Chloranhydrides. — Chromic  acid  forms  chloranhy- 
drides  similar  to  those  of  sulphuric  acid  (p.  195).  Corresponding  to  sul- 
phuryl  chloride,  SO2CI2,  we  have  chromyl  chloride,  Cr02Cl2 ; and  for  the 

r Cl  (Cl 

first  sulphuric  acid  chloranhydride,  SO2  j is  the  salt,  CrOg  j 


VI  r\  VI  r\ 

Cr02<oK 


VI 

CrO, 


,OK 

OK- 


Chromyl  Chloride,  CrOgCla,  Chromhwi  oxychloride,  is  produced  by 
heating  a mixture  of  potassium  (or  sodium)  dichromate  (or  monochro- 
mate) and  sodium  chloride  with  sulphuric  acid  : 

Na2Cr04  -p  2NaCl  -P  4H2SO4  = Cr02Cl2  -p  4NaH804  4-  2H2O, 

and 

Na2Cr207  -p  4NaCl  + 611280^  = 2Cr02Cl2  + 6NaH804  + 3H2O. 

The  water  produced  at  the  same  time  must  be  absorbed  by  the  excess 
of  sulphuric  acid. 

To  prepare  chromyl  chloride,  first  fuse  salt  (10  parts)  with  potassium  dichromate  (12 
parts)  or  with  potassium  monochromate  (17  parts).  The  yellowish-brown  mass  is  broken 
into  coarse  ])ieces,  jdaced  in  a retort  provided  with  a condenser,  and  anhydrous  or  slightly 
fuming  sul])huric  acid  (30  ])arts)  poured  over  them.  When  a gentle  heat  is  applied 
chromyl  chloride  distils  over  aiul  is  purified  by  further  distillation. 

Chromyl  chloride  is  a red,  transparent  liquid,  of  specific  gravity  1.91 
at  25°,  and  fumes  strongly  in  the  air.  It  boils  at  1 16-118°  ; its  vapor 
density  corres])onds  to  the  molecular  formula  Cr0.2Cl2. 


DERIVATIVES  OF  CHROMIC  ACID.  381 

With  water  it  is  decomposed  according  to  the  following  equation; 

CrO.Cl.,  + H2O  = CrOj  + 2HCI. 

Cl 

Chloro-chromic  Acid,  CiO.Xqp^  (see  p.  380),  is  only  known  in 

its  salts.  The  potassium  salt  is  formed  by  heating  potassium  dichromate 
(3  parts)  for  a short  period,  with  concentrated  hydrochloric  acid  (4  parts)  : 

K2Cr207  -f  2HCI  = 2Cr02<^j^  + H2O. 

It  crystallizes  from  the  solution  on  cooling  in  flat,  red  needles.  Heated 
to  100°  it  gives  up  chlorine.  It  is  decomposed  by  water  into  hydro- 
chloric acid  and  potassium  dichromate  : 

2Cr02<Qj^  + H2O  ==  K2Cr207  + 2HCI. 


Bromides  and  iodides  of  chromic  acid  are  not  known,  but  Chromyl  Fluoride,  Cr02Fl2, 
a red-colored  and  very  volatile  liquid,  which  is  decomposed  by  water  with  the  formation 
of  chromic  and  hydrofluoric  acids,  is  obtained  by  heating  a mixture  of  lead  chromate  and 
calcium  fluoride  with  concentrated  sulphuric  acid  in  platinum  or  lead  vessels.  It  etches 
glass.  It  was  formerly  thought  to  be  chromium  hexafluoride,  CrFlg. 


The  following  reaction  is  very  characteristic  for  chromic  acid  : On 
adding  hydrogen  peroxide  to  a solution  of  chromium  trioxide,  or  tlie 
acidified  solution  of  a chromate,  the  red  liquid  is  colored  blue.  On 
shaking  the  blue  solution  with  ether,  the  latter  withdraws  the  blue  com- 
pound which  O.  F.  Wiede  [Ber.  31  (1898),  516]  considers  to  be  the 
anhydride  of  perchromic  acid,  01*209,  together  with  hydrogen  per- 
oxide. This  chemist  succeeded  in  preparing  alkali  salts  of  the  perchromic 
acid  but  always  combined  with  hydrogen  peroxide.  They  are  violet  in 
color  and  are  readily  decomposed ; their  formula  is  probably 

MeO . CrO^  -f  H2O2, 
and  that  of  the  hypothetical  acid : 

CrO^.  OH. 

They  lose  oxygen  and  become  dichromates.  Silver  nitrate  produces  a 
brownish-red  precipitate  of  silver  chromate,  AggCrO^,  with  aqueous 
solutions  of  the  chromates.  This  silver  salt  dissolves  very  readily  in 
ammonia  and  in  the  mineral  acids. 


inor(;anic  chemistry. 


2.  MOLYBDENUM. 

Mo  = 96.0. 

Molybdenum  is  of  comparatively  rare  occurrence  in  nature  ; usually  as  molybdenite 
(MoS^)  andwulfenite  (PbMo04).  Scheele,  in  1778,  obtained  molybdic  acid  from  molyb- 
denite which  had  until  then  been  considered  grapliite.  lljelm,  in  1790,  prepared  molyb- 
denum as  a silver-white  metal,  of  specific  gravity  8.6,  by  igniting  the  chlorides  or  oxides 
in  a stream  of  hydrogen.  It  is  very  hard,  fuses  at  a higher  temperature  than  platinum, 
and  is  as  ductile  and  as  susceptible  of  polish  as  iron.  Like  the  latter  it  takes  up  carbon  at 
1500°  and  becomes  so  hard  that  it  will  scratch  glass.  When  heated  in  the  air  it  oxidizes 
to  molybdenum  trioxide.  It  is  soluble  in  concentrated  sulphuric  and  nitric  acids.  It  is 
also  converted  by  the  latter  into  in.soluble  molybdenum  trioxide,  M0O3. 

Molybdenum  forms  derivatives  of  the  most  different  types,  so  that  it  is  apparently  a 
bivalent,  trivalent,  ([uadrivalent,  etc. , and  even  an  octivalent  metal.  The  vapor  density  of  its 
chloride,  MoClg,  has  been  determined  and  in  it  the  metal  is  positively  quinquivalent.  In 
its  most  stable  oxide,  MoO.^,  it  is  apparently  .sexivalent  and  in  the  stable  sulphide  MoSj 
quadrivalent  ; but  we  are  ignorant  as  to  the  constitution  of  these  bodies  as  well  as  of 
their  molecular  magnitudes.  Only  the  most  important  will  be  described. 

Molybdenum  Dichloride,  MoClj,  resulting  from  the  trichloride,  MoCl,,  when 
heated  in  a stream  of  carbon  dioxide  (together  with  M0CI4),  is  a bright  yellow,  non- 
volatile powder.  It  is  converted  by  potassium  hydroxide  into  the  hydrate,  Mo(OH)2, 
a black  powder. 

The  molecular  formula  of  this  chloride  (and  that  of  the  corresponding  bromide)  must 
apparently  be  trebled — Mo^Clg,  because  dilute  alkalies  change  it  to  chloromolybdic 
hydroxide,  Cl4Mo3(OII).2.  The  latter  behaves  like  a diacid  base  having  the  radical 
(MogCl^)!!.  Concentrated  haloid  acids  convert  it  into  salts  of  the  formula  (MogCl^)!^.^, 
in  which  only  the  two  R-halogen  atoms  are  precipitated  by  silver  nitrate.  In  the 
chloride  (Mo3Cl4)Cl2  therefore  only  two  chlorine  atoms  are  precipitated  as  silver  chloride. 
On  the  other  hand,  chloromolybdic  hydroxide  (a  bright  yellow  amorphous  powder)  is  also 
a feeble  acid.  It  dissolves  without  change  in  dilute  alkalies  and  is  reprecipitated  from 
its  yellow  solution  by  weak  acids. 

The  iron  cyanide  radical  [Fe(CN)g]  and  many  other  radicals  containing  halogens, 
derived  from  cobalt,  chromium  and  platinum,  conduct  themselves  like  (M03CI4).  In 
these  compounds,  the  halogens  belonging  to  the  radicals  are  obscured — protected — from 
the  reagents  with  which  we  are  accustomed  to  see  them  react  in  their  hydrogen  and  other 
derivatives.  This  is  true  also  of  the  other  parts  of  the  radicals,  of  the  molybdenum  in 
(M03CI4),  of  the  iron  in  [Fe(CN)g],  etc.,  etc.  These  metals  can  be  no  longer  recognized 
by  the  reagents  with  which  they  are  ordinarily  detected.  These  new  relations  and  their 
explanation,  by  the  theory  of  electrolytic  dissociation,  will  be  again  considered  in  con- 
nection with  the  iron-cyanogen  compounds. 

Molybdenum  Trichloride,  M0CI3,  produced  by  gentle  heating  (at  250°)  of  molyb- 
denum pentachloride,  M0CI5,  in  a current  of  hydrogen  or  carbon  dioxide,  is  a reddish- 
brown  powder,  resembling  red  phosphorus,  which,  when  strongly  ignited,  yields  a dark- 
blue  vapor.  It  dissolves  with  a beautiful  blue  color  in  concentrated  sulphuric  acid,  upon 
heating,  with  an  emerald-green  color.  Potassium  hydroxide  converts  it  into  the  hydroxide, 
Mo(011)3,  which  forms  .salts  with  acids.  The  ignition  of  the  hydrate  affords  the  black 
oxide,  MogOg.  Strong  heating  of  the  trichloride  in  a current  of  carbon  dioxide  leaves 
molybdenum  dichloride,  MoCl.^,  and  it  sublimes. 

Molybdenum  Tetrachloride,  M0CI4,  is  a brown,  crystalline  powder,  which  appears 
to  break  up  by  eva])oration  into  pentachloride,  M0CI5,  and  trichloride,  M0CI3.  Molyh- 
denimi  disulphide,  MoS^,  is  i)roduced  by  the  ignition  of  the  trisulphide,  M0S3,  away  from 
air.  It  is  a .shining  black  powder,  which  occurs  native  as  molybdenite,  in  hexagonal, 
graphite-like  crystals,  with  a specific  gravity  of  4.6. 

Molybdenum  Pentachloride,  Mod,-,  is  prepared  by  heating  molybdenite  (MoSj), 
or  molybdenum  in  dry  chlorine  gas.  It  is  a metallic,  shining,  black,  crystalline  mass, 
fusing  at  194°  and  distilling  at  268°  ; its  vapor  density  corresponds  to  the  molecular 
formula  MoC^.  It  fumes  and  deliipiesces  in  the  air,  and  dissolves  in  water  with  hissing. 


MOLYBDENUM.  383 

Its  aqueous  solution  has  a brown  color.  It  dissolves  in  absolute  alcohol  and  ether  with 
a dark-green  color. 

The  hexachloride,  MoClg,  is  not  known,  but  the  oxychlo7'ides,  M0OCI4,  and  M0O2CI2, 
are. 

Bromine  forms  perfectly  analogous  compounds  with  molybdenum. 

Molybdenum  Trioxide,  MoO^,  results  on  roasting  metallic  molybdenum  or  the  sul- 
phide in  the  air.  It  is  a white,  amorphous  mass,  which  turns  yellow  on  heating  ; it  fuses 
at  a red  heat  and  then  sublimes  in  brilliant  rhombic  plates  and  needles.  It  is  insoluble  in 
water  and  acids  ; but  dissolves  readily  in  the  alkalies  and  ammonium  hydroxide.  When 
fused  with  the  alkaline  hydroxides  or  carbonates,  salts  are  produced,  partly  derived  from 
the  normal  acid,  H2MoO^,  and  partly  from  the  poly-acids,  and  correspond  to  the  poly- 
chromates : 

K2M0O4,  K2M02OP  K2Mo30jg,  Na2]Mo^Oj3,  KgMo^024,  etc. 

The  ammonhwi  salt,  (NH^)2Mo04,  is  obtained  by  dissolving  the  trioxide  in  concen- 
trated ammonium  hydroxide.  In  the  laboratory  it  serves  as  a reagent  for  phosphoric 
acid.  Alcohol  causes  it  to  separate  out  of  its  solution  in  crystals  ; upon  evaporation, 
however,  the  salt,  (NH^lgMo.Og^  + 4H2O,  crystallizes  out.  Both  salts  are  decomposed 
by  heat,  leaving  molybdenum  trioxide. 

Nitric  acid  gradually  .separates  a yellow  crystalline  powder  from  the  alkali  molybdate.s, 
MoOi^OH)^,  which  is  only  slightly  soluble  in  water.  It  loses  water  on  drying  and 
changes  to  the  normal  hydrate  Molybdic  Acid,  H2MoO^.  By  the  action  of  reducing 
agents,  such  as  zinc,  stannous  chloride,  and  suljrhur  dioxide,  molybdic  acid  solutions 
are  colored  first  blue,  then  green,  then  brownish-red,  and  finally  brown.  This  is  due  to  the 
formation  of  different  lower  oxides.  The  first  blue-colored  reduction  product  is  formed 
by  warming  a molybdic  acid  solution  with  metallic  molybdenum  ; it  has  the  formula 
Mo.^Og  ( = 2M0O3 . Mo02)‘  product,  by  the  action  of  tin  and  hydrochloric 

acid,  is  molybdenum  sesquioxide,  MO2O3.  Its  compounds  are  usually  yellow  in  color. 
Potassium  permanganate  converts  all  these  lower  oxides  into  molybdic  acid. 

Hydrogen  peroxide  converts  molybdic  acid  and  its  salts  into  derivatives  richer  in 
oxygen  ; these  were  supposed  to  be  derived  from  a permolybdic  acid,  HMoO^.  The 
latest  researches  have  demonstrated  that  their  oxygen  content  varies  and  that  they  can- 
not be  referred  to  a common  parent  body.  It  is  thought  that  they  contain,  like  ozone 
and  hydrogen  peroxide,  oxygen  atoms  = ©2  in  union  with  one  another,  therefore  they  have 
been  designated  ozo molybdates  (Muthmann  and  Nagel,  Ber.  31  (1898),  1836). 

Molybdic  acid  can  also  form  poly-acids  with  phosphoric  and  arsenic  acids,  e.  g\, 
H3PO4 . 11M0O3.  These  complex  pho.sphomolybdic  acids  are  distinguished  by  the  fact 
that  they  form  salts  insoluble  in  dilute  acids  with  the  metals  of  the  potassium  group,  with 
ammonia,  with  thallium  and  with  organic  bases.  Sodium  and  lithium  salts  do  not  yield 
precipitates  with  these  acids.  On  adding  a solution  containing  phosphoric  (or  arsenic) 
acid  to  the  nitric  acid  solution  of  ammonium  molybdate,  there  is  produced  a yellow 
crystalline  precipitate  of  ammonium  phospho  molybdate,  (NH^)3PO^  . 11M0O3  6H2O. 

This  reaction  serves  for  the  detection  and  separation  of  phosphoric  acid. 

Molybdenum  Trisulphide,  M0S3,  is  thrown  down  as  a brown  precipitate  from  acidu- 
lated molybdenum  solutions  by  hydrogen  sulphide.  It  dissolves  in  alkaline  sulphides 
forming  sulpho-.salts.  Ignited  away  from  air  it  is  converted  into  black  molybdenum  disul- 
phide, M0S2,  which  occurs  native  as  molybdenite. 

In  addition  to  these  molybdenum  compounds  in  which  the  element  is  sexivalent 
there  is  a 

Molybdenum  Tetrasulphide,  MoS^.  From  this  are  derived  persulphomolybdic 
acid,  H2M0S5,  and  its  salts,  e.  g.,  K2M0S5  and  KHM0S5. 


3^4 


INORGANIC  CHEMISTRY. 


3.  TUNGSTEN. 

W 184. 

Tungsten  is  found  in  nature  in  tlie  tungstates:  as  wolframite,  (Fe,Mn)WO^,  as 
scheelite,  CaWO^,  and  as  stolzite,  PbWO^. 

d'he  metal  is  obtained,  like  molybdenum,  by  the  ignition  of  the  oxides  or  chlorides  in 
a stream  of  hydrogen,  in  the  form  of  a black  powder,  or  in  steel-gray  crystalline  leaflets, 
having  a specific  gravity  of  19.  It  is  very  hard  and  difficultly  fusible.  It  becomes  tri- 
oxide when  ignited  in  the  air. 

Tungsten  forms  the  following  chlorides:  WCl,^,  WCl,,  WCl-  and  WCl^. 

The  Pentachloride,  WCl.,  is  obtained  by  the  distillation  of  WCl^  in  a current  of 
hydrogen  or  carbon  dioxide,  and  consists  of  shining,  black,  needle-like  crystals.  It 
melts  at  248°  and  boils  at  275°,  forming  a greenish-yellow  vapor,  with  the  density  360 
(WClg—  361.25  ; — 32).  It  affords  an  olive-green  solution  and  a blue  oxide,  W^O^, 

with  water.  It  dissolves  with  a deep-blue  color  in  carbon  bisulphide. 

Tungsten  Hexachloride,  WCl,j,  is  produced  when  the  metal  or  a mixture  of  wolf- 
ramite with  carbon  is  heated  in  a current  of  chlorine.  It  forms  a dark-violet,  crystalline 
mass,  which  melts  at  275°  and  boils  at  346°.  The  vapor  density  corresponds  to  the  for- 
mula WClg.  It  dissolves  in  carbon  bisulphide  with  a reddish-brown  color;  it  forms 
tungsten  trioxide,  W().„  with  water.  The  other  chlorides  are  obtained  from  it  by  heating 
them  alone  or  in  a current  of  hydrogen  gas. 

The  Oxychloride,  WOCl^,  consists  of  red  crystals,  fusing  at  210°  and  boiling  at 
229°  ; its  vapor  density  equals  340  (WOCl^  = 341.8;.  The  Dioxychloride,  W02C1.^, 
sublimes  in  bright  yellow,  .shining  leaflets.  They  are  produced  from  the  other  chlorides 
by  the  action  of  water. 

Tungsten  Trioxide,  WO.^,  is  thrown  out  of  the  hot  .solution  of  tungstates  by  nitric 
acid,  as  a yellow  precipitate,  insoluble  in  acids,  but  dissolving  readily  in  potassium  and 
sodium  hydroxides.  Tungstic  Acid,  WO(OH)4,  is,  however,  juecipitated  from  the 
cold  solution,  but  on  standing  over  sulphuric  acid  it  becomes  \V02(011)2  and  at  100° 
passes  into  ditungstic  acid,  H2W20^  :=  W205(0H)2. 

When  tungstic  acid  is  reduced  in  hydrochloric  acid  solution  by  zinc  it  first  becomes 
blue  (formation  of  W'^205),  and  then  brown,  when  the  salt  of  the  dioxide,  WO2,  is  formed. 
Potassium  permanganate  oxidizes  this  to  tungstic  acid. 

The  salts  of  tungstic  acid  are  perfectly  analogous  to  the  molybdates,  and  are  derived 
from  the  normal  acid  or  the  poly-acids.  The  normal  sodium  salt,  Na2W04  -f-  2H2(), 
and  the  so-called  sodium  meta-tungstate,  Na.^W^O^,  -t-  loIL^O,  are  applied  practically. 
Materials  saturated  with  their  solutions  do  not  burst  into  a flame,  but  smoulder  away 
slowly. 

Tungstic  acid  and  the  tungstates  behave  tow'ard  hydrogen  peroxide  just  like  molybdic 
acid. 

The  reduction  of  the  alkaline  tungstates  (by  fu-sion  with  tin,  etc.)  affords  peculiar  com- 
pounds, e.  g.,  or  Nax(  W03)y  ; these  have  various  colors,  possess  metallic  luster, 

and  are  applied  as  tungsten  bronzes. 

Tungstic  acid  also  combines  with  phosphoric  and  arsenic  acids,  forming  derivatives 
analogous  to  those  of  molybdic  acid  with  the  same  acids. 

d'he  metal  is  used  in  the  manufacture  of  tungsten  steel  ; a slight  quantity  of  it  increases 
the  hardness  of  the  latter  very  considerably.  It  was  first  prepared  in  1783  by  J.  and  F. 
d’Elhujar. 


4.  URANIUM. 

U = 239.5. 

In  nalure  it  occurs  chiefly  as  uraninite,  a compound  of  uranic  and  uranous  oxides, 

U0,2.2U()3=  u/v 

d’he  metal,  first  obtained  in  1840  by  Peligot  on  heating  uranous  chloride  with  sodium, 
lias  a silver-while  color  and  a sjiecific  gravity  of  18.7.  When  heated  in  the  air  it  burns  to 
uranous-uraiiic  oxide.  Its  specific  heat  equals  0.0267,  4s  atomic  volume  is  therefore 


URANIUM. 


385 


6.6.  It  melts  at  about  I5cx)°.  'Iliere  are  two  series  of  uranium  compounds.  In  the  one, 
the  metal  is  a tetrad  UX^  ; these  uranous  or  urano-compounds  are  very  unstable,  and  pass 
readily  into  the  uranic  or  derivatives  of  sexivalent  uranium.  Uranous  oxide  is  of  a basic 
nature,  and  only  forms  salts  with  acids. 

'J'he  compounds  of  sexivalent  uranium  are  called  the  uranic  compounds.  The  oxide, 
UC);^,  and  the  hydroxide,  U().^(OH)2,  have  a predominant  basic  character,  but  are  also 
capable  of  forming  salts  witli  bases  which  are  called  loaiiates.  In  the  salts  derived  from 
acids,  e.  g.^  U0.2(N03)2  and  UO^.  SO^,  the  group  UO.^  plays  the  role  of  a metal  ; it  is  called 
tiranyl,  and  its  salts  are  termed  iiranyl  salts.  They  may  also  be  regarded  as  basic  salts. 


URANOUS  COMPOUNDS. 

Uranous  Chloride,  UCI4,  is  obtained  by  heating  metallic  uranium  in  a stream  of  chlo- 
rine, or  uranous  oxide  in  hydrochloric  acid.  It  consists  of  dark-green  octahedra  with 
metallic  luster.  It  volatilizes  at  a red  heat,  forming  a I'ed  vapor,  whose  density  agrees 
with  the  formula  UCI4.  It  deliquesces  in  the  air,  and  dissolves  with  hissing  in  water. 
Uranous  hydroxide  remains  when  the  solution  is  evaporated. 

Uranous  Oxide,  UO2,  is  formed  when  the  other  oxides  are  heated  in  a current  of 
hydrogen.  It  is  a black  powder,  which  becomes  uranous-uranic  oxide,  UO2.  2UO3,  when 
heated  in  the  air. 

Uranous  oxide  dissolves  with  a green  color  in  hydrochloric  and  concentrated  sulphuric 
acids.  Uranous  sulphate,  U(  804)2  + ^^12^5  consists  of  green  crystals.  From  the  salts 
the  alkalies  precipitate  the  voluminous,  bright  green  uranons  hydroxide,  U(OH)4,  which 
becomes  brown  on  exposure.  Uranous  salt  solutions  are  distinguished  by  a very  remark- 
able absorption  spectrum. 


SEXIVALENT  URANIUM  COMPOUNDS. 

Uranium  Hexachloride,  UClg,  has  not  been  obtained,  but  the  oxychloride, 
UO2CI2  (uranyl  chloride),  exists  ; it  is  made  by  heating  UO2  in  dry  chlorine  gas,  or 
by  the  evaporation  of  uranyl  nitrate  with  hydrochloric  acid.  It  is  a yellow  crystalline 
mass,  deliquescing  in  the  air. 

Uranic  Oxide,  UO3  or  Uranyl  Oxide,  UO2.  O,  is  a yellow  powder,  and  is  obtained 
by  heating  uranyl  nitrate  to  250°.  When  warmed  with  nitric  acid  it  becomes  uranyl 
hydrate  or  tira?iic  acid,  U02(OH)2,  which  is  also  yellow-colored. 

Uranyl  Nitrate,  U02(N03)2,  results  from  the  solution  of  uranous  or  uranic  oxide,  or 
more  simply  of  uraninite  in  nitric  acid.  It  crystallizes  with  six  molecules  of  water,  in 
large,  greenish-yellow  prisms,  which  are  readily  soluble  in  water  and  alcohol.  On  add- 
ing sulphuric  acid  to  the  solution,  Uranyl  Sulphate,  UO2SO4  -f  6H2O,  crystallizes  out, 
on  evaporation,  in  lemon-yellow  needles. 

If  sodium  or  pota.ssium  hydroxide  be  added  to  the  solutions  of  uranyl  salts,  yellow  pre- 
cipitates of  the  uranates,  K2U20^  and  Na2U207,  are  obtained.  These  are  soluble  in 
acids.  In  commerce  the  sodium  salt  is  known  as  iiraniiiin  yellow,  and  is  employed  for 
the  yellow  coloration  of  glass  (uranium  glass)  and  porcelain.  The  uranates  can  be 
obtained  in  crystalline  form,  by  igniting  uranyl  chloride  with  alkali  chlorides  in  the  pres- 
ence of  ammonium  chloride.  Nascent  hydrogen  reduces  uranic  to  uranous  compounds. 

The  so-called  uranic-uranous  oxide,  which  constitutes  uraninite,  and  is  formed  by  the 
ignition  of  the  other  oxides  in  the  air,  must  be  viewed  as  uranous  uranate,  2U()3.  UO.,  = 
VI  IV 

(U02.0.2)2U.  _ 

Many  uranium  salts  exhibit  magnificent  fluorescence.  The  oxide  colors  glass  fluxes  a 
beautiful  greenish-yellow  (uranium  gla.ss).  Uranous  oxide,  UO2,  imparts  a beautiful 
black  color  to  glass  and  porcelain. 


33 


386 


INORGANIC  CHEMISTRY. 


Besides  these  compounds,  in  which  uranium  aj)pears  to  be  quadrivalent  and  sexivalent, 
it  also  yields  a pentachloride,  UCI5,  like  molybdenum  and  tungsten.  'I  he  same 
results  on  conducting  chlorine  gas  over  a moderately  heated  mixture  of  carbon  with  one 
of  the  uranium  oxides.  It  consists  of  dark  needles  which,  in  direct  light,  are  metallic 
green,  but  in  transmitted,  ruby-red.  It  delicpiesces  in  the  air  to  a yellowish-green  liquid  ; 
upon  heating  it  is  dissociated  into  uranous  chloride  and  chlorine  (at  120-235°). 

d'here  is  also  a Ictroxide,  U()^,  which,  like  the  trioxide,  UO^,  yields  salts  with  the 
bases.  It  corresponds  to  molybdenum  tetrasulphide  MoS^  (p.  383). 


MANGANESE. 

Mil  = 55.0. 

According  to  its  atomic  quantity,  manganese  bears  the  same  relation  to 
the  elements  of  the  chlorine  group  as  chromium  does  to  the  elements  of  the 
sulphur  group.  The  relationship  manifests  itself  distinctly  in  the  higher 
states  of  oxidation.  Permanganic  acid,  HMnO^,  and  its  anhydride, 
Mn^O^,  are  perfectly  analogous  to  perchloric,  HCIO4,  and  periodic,  HIO4, 
acids  and  their  anhydrides,  CbO^  and  l20^.  The  permanganates  and  the 
perchlorates  are  very  similar,  and  for  the  most  part  are  isomorphous.  The 
manganese  in  them  appears  to  be  septivalent,  like  the  halogens  in  their 
highest  state  of  oxidation.  The  similarity  of  manganese  to  the  halogens 
is  restricted  to  this  one  point  of  resemblance.  In  the  rest  of  its  deriva- 
tives, manganese  shows  great  resemblance  to  the  elements  standing  in  the 
same  horizontal  series  of  the  periodic  system,  viz.,  with  chromium  and 
iron  (p.  393).  Like  these  two  elements,  it  forms  compounds  of  the  types  : 

MO  (M(0H)2.MX2),  M.Pa  (M(On)3.MX3,  MO3  (MOp^). 

1.  In  the  manganous  derivatives,  MnX2,  the  metal  is  bivalent.  These 
salts  are  the  more  stable,  and  comprise  the  most  common  manganese 
compounds.  They  resemble  and  are  usually  isomorphous  with  the  oi4S 
salts  of  iron  and  chromium,  and  the  salts  of  metals  of  the  magnesium 
group  (p.  313). 

2.  The  manganic  compounds,  MnX3,  are  similar  to  and  isomorphous 
with  the  ferric,  chromic  and  aluminium  derivatives;  they  are,  however, 
less  stable,  and  easily  reduced  to  the  manganous  state. 

3.  The  derivatives  of  manganic  acid, 

II^MnO^  = Mn02(0H)2, 

in  which  manganese  is  sexivalent,  are  analogous  to  those  of  ferric 
(IL^FeOJ  and  chromic  acid  (H2C1O4),  and,  of  course,  to  those  of 
sulphuric  acid  (H2SOJ. 

4.  Manganese  also  yields  derivatives  which  can  be  referred  to  the  type 
M.2O7.  These  are  the  perma}iganates,^\\\(^  correspond  in  formula  to  the 
persulphates.  'Phe  manganese  in  them  is  apparently  sei)tivalent,  but  we 
liave  no  definite  knowledge  of  its  structure,  not  even  as  to  whether  it  cor- 
res|)(;nds  to  that  of  the  ])crsul])hates  (]).  297). 

Conseciueiill V,  in  manganese  we  plainly  observe  how  the  similarity  of 
the  elements  in  their  compounds  is  influenced  by  the  valence  (see  p. 


MANGANOUS  COMPOUNDS. 


387 


329).  In  the  ma.nga.uous  condition,  manganese,  like  the  elements  of  the 
magnesium  group,  has  a rather  strong  basic  character,  which  diminishes 
considei'ably  in  the inangan/V  state.  Sexivalent  manganese  has  a metal- 
loidal  acidic  character,  and,  in  manganic  acid,  approaches  sulphur.  By  the 
further  addition  of  oxygen,  manganese  finally  (in  permanganic  acid)  ac- 
quires a metalloidal  character  and  resembles  the  halogens.  We  have 
already  noticed  that  many  other  metals,  especially  chromium  and  iron, 
exhibit  a similar  behavior.  Osmium  in  its  tetroxide,  OsO^,  wholly  resem- 
bles the  halogens. 

On  the  other  hand,  the  metalloidal  and  the  weak  basic  metals  acquire  a strong  basic, 
alkaline  character,  by  the  addition  of  hydrogen,  or  hydrocarbon  groups  (CH.^,  C2Hg). 
The  univalent  groups,  NH^  (ammonium),  P(CH3)^  (tetramethyl  phosphonium), 

(triethyl  sulphine),  Sn(C2H5)3  (tin  triethyl),  etc.,  are  of  metallic  nature,  because  their 
hydroxides,  P(CH3)4.0H,  8(02^13)3.  OH,  80(02113)3 . OH,  are  perfectly  similar  to  the 
hydroxides  (KOH,  NaOH)  of  the  alkali  metals. 


Manganese  is  widely  distributed  in  nature.  It  is  found  native  in 
meteorites.  Its  most  important  ores  are  pyrolusite,  Mn02,  hausmannite, 
Mn304,  braunite,  Mn203,  manganite,  Mn203.H20,  and  rhodochrosite, 
MnC03. 

Metallic  manganese  is  obtained  by  igniting  the  oxides  with  aluminium 
powder  (p.  347).  It  has  a grayish-white  color,  is  very  hard,  and  fuses 
with  difficulty  (at  about  1900°),  and  a little  above  this  temperature  it 
volatilizes.  Its  specific  gravity  is  7. 2-8.0.  It  oxidizes  readily  in  moist 
air.  It  decomposes  water  on  boiling,  and,  when  dissolved  in  acids, 
forms  manganous  salts. 

The  heat  of  formation  of  the  most  important  manganese  compounds  corresponds  to 
the  symbols  : 

(Mn,0,H20)  = 95  (Mn,Cl2)  =^112  (Mn,Cl2,Aq)  = 128 

(Mn,02,H20)  = 1 16  (Mn,S,OJ  = 245  (Mn,0„K)  = 195. 


MANGANOUS  COMPOUNDS. 

Manganous  Oxide,  MnO,  results  from  ignition  of  the  carbonate, 
with  exclusion  of  air,  and  by  heating  all  manganese  oxides  in  hydrogen. 
It  is  a greenish,  amorphous  powder,  which,  in  the  air,  readily  oxidizes 
to  manganous-manganic  oxide,  MiigO^. 

Manganous  Hydroxide,  Mn(OH)2,  is  a voluminous,  white  precipi- 
tate, formed  by  the  alkalies  in  manganous  solutions.  When  exposed  to 
the  air,  it  oxidizes  quickly  to  brown  manganic  hydroxide,  Mn2(OH)3. 

Manganous  salts  usually  have  a ])ale,  reddish  color,  and  are  formed  by 
the  solution  of  manganese  or  oxides  of  manganese  in  acids. 

Manganous  Chloride,  MnCl2,  crystallizes  with  four  molecules  of 
water  in  reddish  tables.  When  heated,  it  decomposes  into  the  oxide, 


388 


INORGANIC  CHEMISIRV. 


liydrochloric  acid  and  water.  Anhydrous  manganous  cldoride  is  ])re- 
pared  by  igniting  the  double  salt 

MnCl,.2Nll,Cl  4 up 

(see  Magnesium  Chloride),  or  by  heating  manganese  oxides  in  hydro- 
chloric acid  gas;  it  is  a crystalline,  reddish  mass,  which  deliquesces  in 
the  air. 

Manganous  Sulphate,  MnSO^,  crystallizes  below  6°  with  seven 
molecules  of  water  (like  magnesium  and  ferrous  sulphates),  at  ordinary 
temperatures  with  sH.p  (like  copper  sulphate);  and  above  25°  with 
qHjO  ; upon  drying,  the  salt 

Mn.SO,  + 11,0 

remains;  the  last  molecule  of  water  does  not  escape  until  at  280°.  It 
forms  double  salts  with  the  alkaline  sulphates,  e.  g.  : 

MnSO^.  K,SO^ -I- 611,0. 

Manganous  Carbonate,  MnCO,,  exists  in  nature  as  rhodochrosite, 
and  is  precipitated  by  alkaline  carbonates  from  manganous  solutions,  as 
a white  powder,  which  turns  brown  on  exposure  to  the  air. 

Manganous  Sulphide,  MiivS,  is  found  in  nature  as  alabandite  or 
manganese-blende.  Alkaline  sulphides  precipitate  a flesh  colored  sul- 
phide from  manganous  solutions.  It  becomes  brown  in  the  air. 


MANGANIC  COMPOUNDS. 

Manganic  Oxide,  Mn203,  manganese  sesquioxide,  is  a brown  or  black 
powder  produced  by  the  ignition  of  all  the  manganese  oxides  in  a current 
of  oxygen  gas.  It  occurs  as  braiinite  in  dark-brown  quadratic  crystals. 

IV  III 

Manganic  Hydroxide,  Mn2(0H)g  or  Mn(OH)3,  manganic  hydrate 
separates  as  a dark-brown  precipitate  on  exposing  ammoniacal  solutions 
of  manganous  salts  to  the  air.  It  dissolves  in  cold  hydrochloric  acid  to 
a dark-brown  liquid,  containing,  in  all  probability,  7fiangani€  chloride, 
MnCl3.  When  this  is  heated  it  decomposes  into  manganous  chloride 
and  chlorine. 

IV 

Manganite,  occurring  in  iron-black  crystals,  is  the  hvdroxide,  Mn202- 

III 

(0H)2  or  MnO.OH. 

Manganous-manganic  Oxide,  MnO.  ^111,03.  consti- 

tutes the  mineral  hausmannite,  crystallized  in  dark-gray  quadratic  octa- 
hedra,  and  is  obtained  as  a reddish-brown  powder  by  the  ignition  of  all 
other  manganese  oxides  in  the  air.  It  reacts  with  hydrochloric  acid, 
according  to  the  equation  : 

MiqO^  + 8IIC1  .iMnCI,  ] 4ll,0  -f-  Cl,. 


MANGANIC  COMPOUNDS. 


389 


Since  manganic  oxide  is  quadratic  in  its  crystallization,  while  all  other  sesquioxides 
(like  corundum  and  hematite)  are  rhombohedral,  and  since  the  first  is  decomposed  by 
dilute  nitric  and  sulphuric  acids  into  manganese  dioxide,  MnOg,  nnd  a manganous  salt, 
it  has  been  generally  supposed  that  manganic  oxide  is  not  a sesquioxide,  but  rather  a 
compound  of  the  dioxide  with  manganous  oxide  : 

MnOj.  MnO  = MnO<(0^]yjj^ 

Hausmannite  is  quadratic,  while  other  metallic  oxides  of  the  type  (the  spinels, 

Y>.  351  and  p.  377,  and  magnetite)  are  isometric;  therefore  the  former  is  not  considered 
a compound  of  manganese  sesquioxide  and  protoxide  ; 

III 

Mn.A-MnO  = “;;0;g>Mn, 
but  as  manganous  oxide  and  the  dioxide  : 

Mn02.2MnO  = MnO<Q-^”>0. 

This  is  shown  by  its  behavior  toward  dilute  nitric  and  sulphuric  acids,  which  decom- 
pose it  into  manganese  dioxide  and  two  molecules  of  manganous  oxide.  Consult  Phanke, 
Jr.  f.  prakt.  Ch.  36  (1887),  451.  Chrysoberyl,  unlike  other  .spinels,  is  trimetric,  and 
other  reactions  clearly  prove  (chiefly  their  deportment  wdth  concentrated  sulphuric  acid) 
that  manganic  and  mangano-manganic  oxides  are  to  be  regarded  as  sesquioxide  derivatives. 


Manganic  oxide,  like  the  other  sesquioxides,  is  a very  feeble  base, 
which  does  not  form  salts  with  dilute  or  weak  acids,  and  by  separation 
of  oxygen  reverts  to  the  manganous  condition.  Its  salts  are  very  un- 
stable. 

Manganic  Sulphate,  Mn2(SOj3,  is  obtained  by  the  solution  of  man- 
ganic oxide,  hydroxide,  or,  better,  manganous-manganic  oxide  in  con- 
centrated sulphuric  acid.  When  the  last  oxide  is  employed  manganous 
sulphate  also  results.  The  best  procedure  is  to  heat  the  hydrate  of  man- 
ganese dioxide  (see  p.  390)  with  concentrated  sulphuric  acid  to  168°, 
when  the  sulphate  will  separate  as  an  amorphous,  dark-green  powder.  It 
dissolves  with  a dark-red  color  in  a little  water.  It  is  said  to  form  alu7ns, 
with  potassium  and  ammonium  sulphates,  e.  g.,  Mn2(SOj3 . K^SO^  -|- 
24H2O,  but  Christensen’s  researches  make  this  rather  doubtful. 


Manganese  Dioxide,  Mn02,  peroxide.  This  is  the  mineral  pyro- 
lusite,  occurring  in  dark-gray  radiating  masses,  or  in  almost  black  rhom- 
bic prisms,  which  possess  metallic  luster.  When  gently  heated  it  is  con- 
verted into  oxide,  by  strong  ignition  into  manganous-manganic  oxide: 

3Mn03  = Mn30,  d-  O2. 

It  is  used  for  making  oxygen.  Chlorine  escapes  when  it  is  warmed 
with  hydrochloric  acid  (p.  49)  : 

Mn02  + 4IICI  = MnClj  + 2II2O  + Cl,. 


390 


INORGANIC  CFIKMISTRY. 


'I'he  dioxide  may  be  obtained  artificially  by  heating  manganous  nitrate 
to  150-160°.  Its  hydrates,  MnO.^.H./)  and  Mn(b.  21 1,0,  are  jirodiired 
on  adding  a hypochlorite  to  the  solution  of  a manganous  salt,  or  if  chlo- 
rine be  conducted  through  a solution  of  manganese  containing  sodium 
carbonate,  or  by  adding  potassium  permanganate  to  a boiling  solution  of 
a manganous  salt,  d'he  precijiitated  dioxide  dissolves  in  cold  hydro- 
chloric acid,  without  liberating  chlorine,  as  manganese  tetrachloride, 
MnCl^,  is  probably  formed;  when  heat  is  applied  it  breaks  down  into 
manganous  chloride,  MnCIa,  and  chlorine,  Clj.  This  deportment  would 
indicate  that  manganese  is  quadrivalent  in  the  dioxide  (p.  259).  Man- 
ganese dioxide  also  unites  with  bases,  yielding  the  so-called  fnangam'tes, 
e.g.,  BaMn^O^and  K.^MiigO^. 

Manganese  ])eroxide  (also  Mn.,03  and  Mn30^)  serves  chiefly  for  the  manufacture  of 
chlorine  gas,  and  it  is,  therefore,  important  from  a technical  ])oint  to  estimate  the  ([uan- 
tity  of  chlorine  which  a given  dioxide  of  manganese  is  able  to  set  free.  This  is  done  by 
boiling  the  oxide  with  hydrochloric  acid,  conducting  the  liberated  chlorine  into  a potas- 
sium iodide  solution,  and  determining  the  separated  equivalent  amount  of  iodine  by  means 
of  sodium  hyposulphite.  Or  the  oxide  is  heated  in  a flask  with  oxalic  and  sulidiuric 
acids,  when  the  oxalic  acid  is  oxidized  to  carbon  dioxide,  and  from  the  quantity  of  this 
set  free  we  can  calculate  the  quantity  of  active  or  available  oxygen  in  the  manganese  oxide  : 

MnO^  + H2C,0,  + H.,SO^  = MnSO^  + 2II2O  d-  2CO.,. 

In  the  preparation  of  chlorine  the  manganese  is  found  in  the  residue  as  manganous 
chloride.  With  the  relatively  high  value  of  pyrolusite,  it  is  important  for  trade  that  the 
peroxide  be  recovered  from  the  residue.  This  regeneration  is  at  present  largely  executed 
by  the  method  proposed  by  Weldon,  according  to  which  the  manganous  chloride,  contain- 
ing an  excess  of  hydrochloric  acid,  is  neutralized  with  lime,  the  clear  liquid  brought  into 
a tall  iron  cylinder  (the  oxidizer),  milk  of  lime  added  and  air  forced  in.  The  mixture 
becomes  warm,  and  so-called  calcium  manganite,  CaMnOg  = CaO . Mn02,  is  precipi- 
tated as  a black  mud  (Weldon’s  mud)  : 

MnCb  -f  2CaO  -f  O = CaMnOg  + CaCb,. 

The  calcium  chloride  solution  is  run  off,  and  the  residual  calcium  manganite  employed 
for  the  preparation  of  chlorine,  when  it  conducts  itself  as  a mixture  of  Mn02  -f  CaO. 


COMPOUNDS  OF  MANGANIC  AND  PERMANGANIC  ACIDS. 

When  oxygen  compounds  of  manganese  are  heated  in  the  air  in  con- 
tact with  potassium  hydroxide,  or,  better,  with  oxidizing  substances,  like 
niter  or  potassium  chlorate,  a dark-green  amorphous  mass  is  produced, 
which  dissolves  in  cold  water  with  a dark-green  color.  When  this  solu- 
tion is  evaporated  under  the  air-pump,  dark-green,  metallic,  rhombic 
jirisms  of  potassium  manganate,  K^MnO^,  crystallize  out.  This  salt 
is  isomorphous  with  potassium  sulphate  and  potassium  chromate.  It  suf- 
fers no  change  by  solution  in  potassium  or  sodium  hydroxide,  but  is  de- 
comj^osed  by  water,  brown  hydrated  manganese  dioxide  separating,  and 
the  green  solution  of  the  manganate  changing  into  a dark-red  solution  of 
the  jiermanganate,  KMnO^: 

3K2Mnf)^  y 3H2O  = 2KMnO^  -j-  Mn02.H2C  -f  4KOII. 


POTASSIUM  PERMANGANATE. 


391 


A similar  conversion  of  the  green  manganate  into  red  permanganate 
occurs  more  rapidly  under  the  influence  of  acids  or  chlorine : 

3K2MnO^  + 4HNO3  = 2KMn04  + MnO^  + 4KNO3  + 2H3O  ; 

K^MnO^  + Cl  = KMnO^  -f  KCl. 

The  red  solution  again  assumes  a green  color  on  the  addition  of  a con- 
centrated solution  of  an  alkaline  hydroxide. 

Owing  to  this  ready  alteration  in  color  the  solution  of  the  manganate 
is  called  chameleon  77imeraL 

Potassium  Permanganate,  KMnO^,  is  best  prepared  by  conducting 
carbon  dioxide  into  the  manganate  solution  until  the  green  color  has 
passed  into  a red.  Hydrated  peroxide  of  manganese  is  i)recipitated. 
When  the  solution  is  concentrated  the  salt  crystallizes  in  dark-red 
rhombic  prisms  isomorphous  with  potassium  perchlorate,  KCIO^.  It  is 
soluble  in  twelve  parts  of  water  at  ordinary  temperatures. 

The  permanganate  solution  is  a strong  oxidizing  agent.  If  the  oxida- 
tion takes  place  in  the  presence  of  a sufficient  quantity  of  acid  the  per- 
manganate is  reduced  to  a faintly  colored  manganous  salt : 

2KMnO,  + 3H2SO,  = 2MnSO,  + K^SO^  + 3H,0  + 5O. 

When  a permanganate  solution  is  added  to  an  acidulated  ferrous  solution, 
the  former  is  decolorized,  and  there  results  a faintly  yellow-colored  solu- 
tion of  ferric  and  manganous  salts: 

2KMn04  + loFeSO^  + SH^SO^  = 2MnSO^  + 5Fe2(SOd3  + SHp  + K.^SO^. 

Hence  the  solution  of  this  salt  serves  for  the  volumetric  estimation  of 
ferrous  salts. 

In  neyfral  ox  alkaline the  permanganate  in  oxidizing  is  decom- 
posed with  the  separation  of  manganese  oxides  : 

2KMn04  -f  HjO  = 2Mn02  -f-  2KOH  -|-  3O. 

In  the  same  manner,  the  permanganate  oxidizes  and  destroys  many 
organic  substances,  therefore  its  solution  cannot  be  filtered  through  paper 
without  decomposition.  It  serves  as  a disinfectant. 

The  permanganate  is  also  reduced  by  hydrogen  peroxide  in  acid  solu- 
tion (p.  103). 

The  remaining  permanganates  are  similar  to  and  isomorphous  with  the 
perchlorates.  The  sodium  salt  is  very  soluble  in  water,  and  does  not 
crystallize  well. 

Very  cold  sulphuric  acid  added  to  dry  permanganate  causes  the 
separation  of  Manganese  Heptoxide,  Mn^O^,  an  oily,  dark-colored 
liquid.  By  careful  warming  it  is  converted  into  dark-violet  vapors, 
which  explode  when  heated  rapidly.  Manganese  heptoxide  has  a 
violent  oxidizing  action ; ])aper,  alcohol  and  other  organic  matter  are 
inflamed  by  mere  contact  with  it. 


392 


inor(;anic  cukmis'ikv. 


METALS  OF  GROUF"  VIII. 

The  last  group  of  the  periodic  system  comprises  the  elements: 

Fe  — 56  Ni  = 58.7  Co  = 59 

Ru  — 101.7  Kh  = 103.0  Pel  = 106 

Os  — 191  Ir  193-0  Pt  = 194.8. 

These  elements  are  the  middle  members  of  the  three  great  periods,  and 
they  have  no  analogues  in  the  two  small  jjeriods  (]).  245). 

As  regards  both  atomic  weights  and  i)hysical  and  chemical  dei)ortment, 
these  elements  constitute  a transition  from  the  ])receding  members  of  the 
great  periods  (Cr  and  Mn  ; Mo;  W)  to  the  next  following  members  (Cu, 
Ag,  An,  and  Zn,  Cd,  Hg).  The  elements  standing  side  by  side  (heterol- 
ogous) and  belonging  to  the  same  periods  are  very  similar  in  their 
physical  properties,  and  show,  e.  g.,  very  close  si)ecific  gravities.  They 
are,  therefore,  usually  arranged  in  groups,  and  are  distinguished  as  (i)  the 
iron  group  (Fe,  Ni,  Co),  with  the  specific  gravity  7. 8-8. 6;  (2)  thegrouj) 
of  the  light  platinum  metals  (Ru,  Rh,  Pd),  with  the  specific  gravity 
11.8-12. 1 ; and  (3)  the  group  of  the  heavy  platinum  metals  (Os,  Ir,  Ptj, 
with  the  specific  gravity  21. 1-22.4. 

On  the  other  hand,  the  homologous  elements  (Fe,  Ru,  Os;  Ni,  Rh, 
Ir ; and  Co,  Pd,  Pt)  show  a like  similaritv  in  their  chemical  i)roperties, 
as  do  the  other  homologous  elements.  This  re.semblance  shows  itself 
chiefly  in  their  combination  forms,  and,  of  course,  too,  in  the  properties 
of  the  compounds  (p.  329).  We  know  that  the  metals  of  group  VI 
(chromium,  molybdenum,  tungsten)  and  of  group  VII  (manganese)  form 
the  highest  oxides  (MeOg  and  Me20^)  having  an  acidic  nature.  In  the 
adjacent  elements  of  group  VIII  (iron,  ruthenium,  and  osmium)  we  find 
salts : 

KgFeO^,  K2RUO4,  K20s0^, 

derived  from  the  unstable  trioxides  FeOg,  RuOg  and  OsOg.  This  acid- 
forming function  disappears  in  the  following  members,  Ni,  Rh,  Ir,  and 
Co,  Pd,  Pt ; their  chemical  valences  diminish  rapidly  and  they  attach 
themselves  to  Cu,  Ag  and  Au. 

Consequently  the  whole  physical  and  chemical  deportment  of  the  nine 
elements  about  to  be  considered  is  governed  by  their  position  in  the 
periodic  system. 


METALS  OF  THE  IRON  GROUP. 

'I'he  metals  of  this  group,  iron,  nickel  and  cobalt,  are  distinguished 
from  all  the  other  elements  by  their  magnetic  ])ro])erties. 

Iron  forms  three  series  of  compounds  after  the  forms,  FeOg,  Fe20g  and 
lAO.  In  its  highest  combinations  iron  has  an  acidic  character,  and  the 
derivatives  of  ferric  acid  (H2FeOJ  are  perfectly  similar  to  those  of 


IRON. 


393 


chromic  and  manganic  acids;  they  are,  however,  less  stable  than  the 
latter.  Their  analogues  with  cobalt  and  nickel  are  unknown. 

The  ferrzV  compounds,  FeXg,  are  much  like  the  aluminium,  chromic 
and  manganic  derivatives.  They  are  generally  isomorphous  with  them. 
They  are  characterized  among  iron  salts  by  their  relative  stability.  The 
highest  oxides  of  cobalt  are  far  less  stable,  while  the  higher  salts  with 
nickel  are  unknown. 

Again,  iron,  nickel  and  cobalt  form  ous  compounds  (FeX^,  NiX^, 
C0X2)  in  which  they  appear  to  be  dyads.  They  resemble  the  compounds 
of  chromium,  manganese,  and  copper  of  the  same  form,  and  those  of  the 
magnesium  metals.  The  ferrous  salts  are  not  as  stable  as  the  ferric;  they 
are  readily  oxidized  to  the  latter. 

The  cobaltous  and  nickelous  compounds  are  quite  stable,  and  in  this 
respect  these  metals  ally  themselves  with  co})per  and  zinc. 


1.  IRON. 

Fe  = 56.0. 

This  very  important  metal  is  widely  distributed  in  nature.  It  is  found 
native  on  the  earth’s  surface  almost  exclusively  in  meteorites,  which 
generally  contain  nickel.  Iron  granules  free  from  the  last  metal  are  said 
to  occur  in  the  coal  measures  of  Missouri.  Iron  is,  however,  present  in 
great  masses  in  other  worlds  which  (like  the  sun)  are  surrounded  by  an 
atmosphere  of  glowing  hydrogen,  the  heat  of  which  is  so  intense  that 
only  the  free  elements  can  exist  side  by  side  (p.  29). 

The  most  important  iron  ores  are : magnetite  (FegOJ,  hematite 
(Fe203),  brown  iron  ore  or  limonite  (hydrated  oxide)  and  siderite 
(FeCOg).  These  ores  constitute  almost  the  sole  material  for  the  manu- 
facture of  iron  ; the  sulphide  ores,  like  pyrite,  are  not  so  well  adapted 
for  iron  making. 

When  iron  is  mentioned  a distinction  must  be  made  between  the 
metal  iron,  the  pure  element,  and  that  iron  which  meets  with  a universal 
application  in  the  arts  and  industries.  The  latter  is  really  not  a metal 
but  an  alloy  of  iron  with  carbon,  and  frequently  also  with  silicon,  phos- 
phorus, sulphur,  and  manganese.  Chemically  pure  iron,  the  real  metal 
iron,  is  not  applied  technically  because  its  production  would  be  too  ex- 
pensive. Hence,  a distinction  must  be  made  between  chemically  pure 
iron  and  the  technical  iron. 

A.  Chemically  pure  iron. 

Chemically  pure  iron  is  obtained  by  heating  the  pure  oxide  or  the  oxa- 
late in  a current  of  hydrogen  : 


Fe^Og  -f  3II2  = 2Fe  -p  ; 

to  have  complete  reduction  the  temperature  finally  should  reach  600°. 
If  the  reduction  occurs  at  a red  heat,  the  powder  glows  in  the  air,  and 
burns  (pyrophoric  iron).  The  strongly  ignited  powder  is  not  inflam- 


394 


inoR(;anic  cukmistry. 


nuible.  Iron  obtained  by  the  ele'clrolysis  of  ferrous  sulj)hate  always  ccni- 
lains  some  hydrogen. 

Chemically  pure  iron  when  hammered  has  a silver-white  color,  is  toler- 
ably soft,  and  is  magnetizable.  Its  specific  gravity  is  7.78.  It  melts 
transitionally  at  1800°  and  vaporizes  at  higher  temjieratures.  In  dry  air 
it  alters  slowly  but  rusts  rapidly  in  moist  air.  Boiling  water  is  decom- 
posed by  finely  divided  ])ure  iron  with  evolution  of  hydrogen.  It  dis- 
solves readily  in  acids,  and  in  nitricacid  with  the  liberation  of  nitric  oxide. 

B.  Technical  Irofi.  d'his  always  contains  carbon,  d’he  carlxm  con- 
tent increases  as  the  temj^erature,  at  which  the  ore  is  fused  with  the  coal, 
rises.  However,  even  at  the  most  elevated  temiieratures  iron  does  not  take 
up  more  than  5 per  cent,  of  carbon.  The  latter  is  either  mechanically 
mixed  as  graj)hite  with  the  iron  or  it  exists  as  iron  carbide,  he„C,„  (<?.  g., 
FegC)  chemically  combined — alloyed. 

The  quantity  of  carbon  and  the  form  in  which  it  exists  leads  to  the 
varieties  cast  iron,  steel  and  wrought  iron,  with  their  subdivisions. 

I.  Cast  Iron  contains  2.3  or  more  per  cent,  of  carbon  ; its  melting 
temperature  rises  as  the  carbon  content  falls  (1050-1250°).  It  is  brittle 
and  cannot  be  welded  or  forged.  The  varieties  of  cast  iron  are  : 

{ai)  White  Iron.  All  of  its  carbon  is  chemically  combined.  It  is 
hard,  brittle,  and  white  in  color.  It  is  only  made  for  conversion  into 
wrought  iron.  White  iron  is  distinguished  as  fibrous  and  “ spiegeleisen," 
according  to  the  appearance  of  its  fracture. 

{b)  Gray  Iron.  Only  a portion  of  its  carbon  is  chemically  combined 
with  the  iron.  The  greater  portion  exists  as  graphite.  It  is  softer,  more 
fusible  than  the  preceding  and  of  a bright  gray  to  black  color.  It  is 
used  in  the  manufacture  of  castings,  and  of  malleable  iron ; particularly 
by  the  Bessemer  process. 

In  the  production  of  the  different  kinds  of  iron  not  only  carbon  but 
also  manganese  and  silicon  play  an  important  part.  Their  presence  or 
absence  also  determines  the  nature  of  the  iron.  Manganese  favors  the 
chemical  union  of  iron  and  carbon.  Hence  white  iron  is  usually  man- 
ganiferous.  Silicon  decomposes  the  alloy  of  iron  and  carbon  and  favors 
the  production  of  grajihite.  The  color  of  gray  cast  iron  is  determined 
by  the  size  of  the  graphite  plates.  If  they  are  very  small  the  iron  is  fine- 
grained and  its  color  light;  otherwise  the  aggregation  is  coarsely  crystal- 
line and  the  color  dark — light  and  dark  gray,  and  black  pig  iron. 

H.  Wrought  Iron  and  Steel  with  1.6  per  cent,  and  less  of  carbon. 
(Iron  containing  1.6  to  2.3  ])er  cent,  of  carbon  is  no  longer  manufactured 
technically.)  It  fuses  with  difficulty,  but  is  malleable  and  can  be  welded. 
It  melts  at  1400°  and  higher.  The  method  of  preparation  and  the  con- 
dition in  which  the  malleable  iron  exists  at  the  conclusion  of  its  manu- 
facture distinguish  it  as 

A.  Fluid  Iron. 

B.  Weld  Iron. 

'Fhe  first  has  been  melted  during  its  manufacture  and  is  therefore  free 
from  slag,  while  the  second  varietv  has  only  been  reduced  to  a ]xisty  form 
and  therefore  contains  slag.  When  a glowing  piece  of  weld  iron  is 


IRON. 


395 


suddenly  cooled  {e.  g.,  by  immersion  in  water)  its  hardness  is  appreciably 
increased,  in  which  case  it  is  called  steel,  otherwise  malleable  or  wrought 
iron.  The  following  subdivisions  are  also  known  : 

{tempered  : Fluid  steel, 
not  perceptibly 

tempered  : Fluid  weld  iron. 

{tempered : Weld  steel, 
not  perceptibly 

tempered  : Wrought  iron. 

The  hardening  is  generally  influenced  by  the  carbon  content ; steel 
contains  from  0.6-1. 6 per  cent,  and  wrought  iron  less  than  0.6  per  cent, 
of  carbon.  However,  the  formerly  customary  distinction  between  steel 
and  wrought  iron  on  the  basis  of  carbon  content  alone  is  no  longer  ten- 
able, as  at  present  tempered  varieties  of  iron  are  produced,  containing 
very  considerable  quantities  of  manganese,  nickel,  silicon,  tungsten,  and 
chromium,  and  very  little  carbon. 

The  classification  just  given  of  the  technically  available  kinds  of  iron  was  prepared  by 
an  international  commission  at  the  Centennial  Exposition  in  Philadelphia  (1876),  but  it 
unfortunately  has  not  been  generally  adopted. 

Pig  iron  (crude  iron)  is  always  produced  in  blast  furnaces  by  the  reduc- 
tion of  its  ores;  while  wrought  iron  is  rarely  made  from  ores  directly,  and 
never  in  blast  furnaces.  As  a rule,  wrought  iron  is  manufactured  from  pig 
iron,  as  its  name  would  imply. 

Metallurgy  of  Iron. — The  extraction  of  iron  from  its  oxygen  ores  is 
based  upon  the  reduction  of  the  same  by  carbon  at  a red  heat.  In  the 
oldest  method,  the  ores  were  heated  with  carbon  in  forges;  in  this  way 
the  excess  of  air  consumed  the  greater  portion  of  the  carbon,  and  the  prod- 
uct was  an  iron  poor  in  carbon,  wrought  iron,  a spongy  mass,  which  was 
then  worked  under  the  hammer.  The  present  methods  have  been  de- 
veloped since  the  beginning  of  the  fifteenth  century.  According  to 
these  pig  iron  is  first  prepared  from  the  ores,  and  this  afterwards  con- 
verted into  steel  or  wrought  iron.  The  smelting  of  the  ores  is  executed 
in  large,  walled  blast  furnaces,  that  permit  the  process  to  proceed  without 
interruption.  The  furnaces  are  filled  from  openings  above,  with  alter- 
nating layers  of  coal,  broken  ore  and  fluxes  containing  silica  and  lime; 
the  fluxes  facilitate  the  melting  together  of  the  reduced  iron.  The  air 
necessary  for  the  process  is  blown  into  the  lower  contracted  portion  of 
the  furnace  by  means  of  a blowing  engine.  The  combustion  of  the  coal 
produces  carbon  monoxide,  which  reduces  the  iron  oxides  to  metal : 

FcjOg  3CO  = 2Fe  -f-  3CO2. 

As  the  reduced  iron  sinks  in  the  furnace  it  comes  in  contact  with  the 
coal,  takes  up  carbon  and  forms  cast  iron,  which  fuses  as  it  sinks  lower 
and  collects  in  the  hearth  of  the  furnace.  Protracted  and  strong  heating 
converts  the  chemically  combined  carbon  into  tlie  graphitic  form,  and  thus 
accelerates,  with  the  assistance  of  silicon,  the  formation  of  the  gray  cast 


39<> 


INORGANIC  CHEMISIRY. 


iron.  The  eanhy  impurities  of  tlie  ores  coml)inc  willi  the  fiuxcs  to  a 
readily  fusible  s/dg,  winch  euvel()i)S  the  fused  iron  aud  protects  it  from 
oxidation. 

'I'o  convert  the  cast  iron  thus  produced  into  wrought  iron,  carl^on  must 
be  withdrawn  from  it.  lu  making  the  wrought  iron  the  cast  iron  is 
fused  in  open  hearths  (refining  process),  or  in  reverberatory  furnaces 
with  air  access,  aud  the  mass  stirred  thoroughly  until  it  has  become  i)asty 
(puddling  process).  In  this  way  almost  all  the  carbon  is  burned  to  car- 
bon monoxide  and  the  other  admixtures,  like  silicon,  sulphur,  and  phos- 
phorus,  present  in  small  quantities,  are  oxidized.  The  wrought  iron  is 
then  worked  up  by  rolling,  or  under  the  iron  hammers  (bar  iron).  If 
the  decarburization  be  not  carried  (piite  as  far  the  j)roduct  will  be  steel 
(puddle  steel).  Such  products  are  not  free  from  slag.  Cast  iron  can  be 
changed  into  malleable  iron  by  annealing.  Castings  of  white  iron  are 
imbedded  in  ferricoxide  powder  and  ex])osed  for  some  time  to  a red  heat. 
A portion  of  the  carbon  is  lost  and  malleable  iron  results. 

Steel  can  be  manufactured  from  wrought-iron  by  cementation.  The 
iron  bars,  mixed  with  fine  charcoal,  are  exposed  to  a red  heat,  when  the 
iron  takes  up  carbon  from  the  surface  inward,  d’he  bars  are  then  re- 
forged, again  heated  with  fine  charcoal,  and  the  proce.ss  repeated  until 
the  mass  becomes  as  homogeneous  as  po.ssible  (cementation  steel).  Cem- 
entation is  just  the  opposite  of  annealing.  A more  homogeneous  steel  is 
obtained  by  refusion  in  crucibles  (cast  steel). 

At  present,  steel  is  chiefly  prejmred  directly  from  pig  iron,  by  the 
method  invented  by  Bessemer,  in  1855.  It  consists  in  blowing  air,  under 
high  pressure,  into  the  molten  iron,  until  the  necessary  amount  of  carbon 
has  been  burned  out  (Bessemer  steel). 

An  iron  rich  in  silicon  (1.5-2  per  cent,  silicon)  is  well  adapted  for  the 
purpose,  because  by  the  simultaneous  combustion  of  the  silicon  the  tem- 
perature is  considerably  increased.  The  operation  is  conducted  in  a 
pear-shaped  vessel  known  as  the  converter.  The  air  is  blown  in  through 
openings  in  the  bottom.  The  decarburization  by  the  Swedish  method 
only  continues  to  the  point  where  steel  is  fused.  This  is  better  accom- 
]dished  by  the  English  method,  which  removes  the  carbon  so  as  to  con- 
vert the  mass  into  wrought  iron,  and  then  it  is  again  carbonized  by  add- 
ing molten  spiegeleisen,  or  by  the  addition  of  coke. 

d’he  Bessemer  process  originally  was  only  adapted  to  crude  iron  con- 
taining as  little  sul]:)hur  and  phosphorus  as  possible  (at  the  highest, 
o 05  per  cent,  phosphorus),  because  in  this  process  the  phosphorus  is  not 
consumed,  but  remains  unaltered  in  the  steel.  By  a slight,  yet  very 
essential  alteration,  Snelus  (1872)  and  Thomas  and  Gilchrist  (1878)  ren- 
dered it  suital)le  for  iron  containing  much  phosphorus.  Their  process  is 
now  known  as  the  ‘‘  basic  process,”  and  consists  in  lining  the  converter 
with  a basic  lining  material  of  burnt  dolomite.  By  contact  with  this  sub- 
stance the  iron  is  comjiletely  de])hosphorized  and  the  jfiiosphorus  changed 
to  calcium  and  magnesium  phospliates. 

'I'lic  basic  (lopliospliorizing  inctbod  lias  aUained  very  great  importance  because  all  the 
jihosphorus  contained  in  iron  ores  collects  in  the  slag  of  the  converter,  d'he  latter  con- 
tains as  much  as  15-20  per  cent,  of  phosphoric  acid  (and  may  even  be  increased  to  24 


FERROUS  COMPOUNDS.  397 

per  cent.),  existing  as  calcium  phosphate,  and  this  may  be  applied  as  a fertilizer  in  agri- 
culture (Thomas  slag)  (p.  305). 

A third  method  (Siemens-Martin)  of  making  steel  consists  in  puddling 
the  different  varieties  of  iron  together,  or  with  iron  ores.  Siemens-Mar- 
tin steel  is  obtained  by  fusing  pig  iron  with  wrought  iron  (old  iron  rails). 
It  is  much  used.  Uchatius  steel  is  prepared  by  fusing  granulated  pig 
iron  together  with  some  iron  ore  and  pyrolusite  in  graphite  crucibles. 


Ordinary  iron,  even  the  purest  wire,  always  contains  foreign  ingredi- 
ents, principally  carbon  and  manganese,  and  minute  quantities  of  silicon, 
sulphur,  phosphorus,  nitrogen,  nickel,  cobalt,  titanium  and  other  metals. 
The  quantity  of  manganese  is  purposely  increased  (to  33  per  cent.),  as  by 
this  means  the  iron  acquires  valuable  technical  properties;  it  becomes 
more  compact  and  solid.  When  iron,  containing  carbon,  is  dissolved  in 
hydrochloric  acid  the  chemically  combined  carbon  unites  with  hydrogen, 
forming  volatile  hydrocarbons,  while  the  mechanically  admixed  graphite 
remains  behind.  The  total  carbon  is  determined  by  the  solution  of  the 
iron  in  bromine  water  or  cupric  chloride,  when  all  the  carbon  remains 
behind. 

Ordinary  iron  rusts  rapidly  in  moist  air,  or  it  becomes  covered  with  a 
layer  of  ferric  hydroxide.  When  ignited  in  the  air  it  is  coated  with  a 
layer  of  ferrous-ferric  oxide  (Fe^OJ  which  is  readily  detached.  It  burns 
with  an  intense  light  in  oxygen. 

In  contact  with  a magnet  iron  becomes  magnetic ; steel  alone  retains' 
the  magnetism,  while  cast  iron  and  wrought  iron  soon  lose  this  property 
after  the  removal  of  the  magnet. 

Iron  decomposes  water  at  a red  heat,  with  the  formation  of  ferrous- 
ferric  oxide,  and  the  liberation  of  hydrogen. 

The  metal  dissolves  readily  in  hydrochloric  and  sulphuric  acids,  with 
evolution  of  hydrogen  ; the  latter  has  a peculiar  odor,  due  to  hydrocar- 
bons which  are  liberated  at  the  same  time.  Iron  dissolves  in  nitric  acid 
with  evolution  of  nitrogen  oxides.  On  dipping  iron  into  concentrated 
nitric  acid,  and  then  washing  it  with  water,  it  is  no  longer  soluble  in 
the  acid  (passive  iron)  ; this  phenomenon  is  probabjy  due  to  the  produc- 
tion of  ferrous  oxide  upon  its  surface. 


FERROUS  COMPOUNDS. 

These  are  produced  by  the  solution  of  iron  in  acids,  and  may  also  be 
obtained  by  the  reduction  of  ferric  salts  : 

Fe2Ck  -h  Zn  = 2FeCl2  -f  ZnCl2. 

In  the  hydrous  state  they  are  usually  of  a green  color;  in  the  air  they 
oxidize  to  ferric  salts: 

2FeO  -|-  O = FcjOj. 


398 


INORGANIC  CHP:MISTKY. 


Ferrous  Chloride,  FcCl^^,  crystallizes  from  aqueous  solutions  in  green 
monoclinic  prisms,  with  four  molecules  of  water.  These  delicjuesce  in 
the  air  and  oxidize.  Particularly  beautiful  crystals  are  obtained  by 
exposing  an  alcoholic  solution  of  ferric  chloride  in  a closed  vessel  to  the 
action  of  direct  sunlight.  The  alcohol  acts  as  a reducing  agent  and  is 
oxidized  to  aldehyde  (Organic  Chemistry).  The  anhydrous  salt  is 
formed  by  conducting  hydrogen  chloride  over  heated  iron.  It  is  a white 
mass,  which  fuses  on  application  of  heat  and  sublimes  at  a red  heat  in 
white,  six-sided  leaflets.  Its  vapor  density  at  1300-1500°  corresponds  to 
the  formula  FeCl.^,  but  it  ajqiears  that  at  lower  temperatures  it  is  also 
possible  for  the  molecules  Fc2Cl^  to  exist. 

It  forms  double  salts  with  the  alkaline  chlorides,  e.  g.  : 

YeC\.2KC\-{-2\\^0. 

Ferrous  Iodide,  Fel2,  is  obtained  by  warming  iron  with  iodine  and 
water.  It  crystallizes  with  four  molecules  of  water. 

Ferrous  Oxide,  FeO,  is  a black  powder,  resulting  from  the  reduction 
of  ferric  oxide  by  carbon  monoxide.  When  warmed  in  the  air  it  oxidizes 
readily.  Ferrous  Hydroxide,  Fe(OH)2,  is  thrown  out  of  ferrous  solu- 
tions by  the  alkalies,  as  a greenish-white  precipitate.  Exposed  to  the 
air,  it  oxidizes,  becoming  reddish-brown.  It  is  somewhat  soluble  in 
water,  and  has  an  alkaline  reaction. 

Ferrous  Sulphate,  FeSO^,  crystallizes  with  seven  molecules  of  water 
in  large,  greenish,  monoclinic  prisms,  and  is  generally  called 
The  crystals  effloresce  somewhat  in  dry  air.  They  oxidize  in  moist  air, 
and  become  coated  with  a brown  layer  of  basic  ferric  sulphate.  At  100° 
they  lose  six  molecules  of  water,  and  change  to  a white  powder.  The 
last  molecule  of  water  escapes  at  300°.  Therefore,  ferrous  sulphate 
behaves  just  like  the  sulphates  of  the  metals  of  the  magnesium  group 
(p.  315).  Like  them,  it  unites  with  alkaline  sulphates  to  double  sulphates, 
which  contain  six  molecules  of  water,  e.  g.,  FeSO^.  K2SO4 -f- 6H2O. 
These  are  more  stable  than  ferrous  sulphate,  do  not  effloresce,  and  oxidize 
very  slowly  in  the  air. 

The  salt  FeSO^.  (NHJ2SO4  -f-  6H2O — Mohr’ s salt — is  particularly  char- 
acterized by  its  stability  in  the  air,  and  is  therefore  employed  to  determine 
the  strength  of  the  potassium  permanganate  used  in  titrations. 

Ferrous  sulphate  is  obtained  by  dissolving  iron  in  dilute  sulphuric  acid. 
Commercially,  it  may  also  be  made  from  pyrite  (FeS2),  which  by  careful 
roasting  loses  one  atom  of  sulphur,  and  is  converted  into  ferrous  sul- 
])hide  (FeS),  which,  in  the  presence  of  water,  absorbs  oxygen  from  the 
air,  and  is  converted  into  sulphate,  which  may  then  be  extracted  by  water. 

Iron  vitriol  has  an  extended  practical  application;  among  other  uses, 
it  is  emj)loyed  in  the  ])reparation  of  ink,  and  in  dyeing. 

When  heated  it  decomposes  according  to  the  following  equation : 

2FeS(),  = Fe2^;t  I L 

On  this  is  based  the  production  of  fuming  Nordhausen  sulphuric  acid 
(p.  194),  and  of  colcothar. 


FERRIC  COMPOUNDS. 


399 


Ferrous  Carbonate,  FeCOg,  exists  in  nature  as  siderite,  crystallized 
in  yellow-colored  rhombohedra,  isomorphous  with  calcite  and  smithsonite. 
Sodium  carbonate  added  to  ferrous  solutions  precipitates  a white  volum- 
inous carbonate,  which  rapidly  oxidizes  in  the  air  to  ferric  hydroxide. 
Ferrous  carbonate  is  somewhat  soluble  in  water  containing  carbon  dioxide, 
hence  present  in  many  natural  waters. 

Ferrous  Phosphate,  FegCPOJg  + SH^O,  occurs  crystallized  in  bluish 
monoclinic  prisms  as  Vivianite.  Precipitated  by  sodium  phosphate  from 
ferrous  solutions,  it  is  a white  amorphous  powder,  which  oxidizes  in  the 
air. 

Ferrous  Sulphide,  FeS,  is  a dark-gray,  metallic  mass,  obtained  by 
fusing  together  iron  and  sulphur.  It  is  made  use  of  in  laboratories  for 
the  preparation  of  hydrogen  sulphide.  If  an  intimate  mixture  of  iron 
filings  and  sulphur  be  moistened  with  water,  the  union  will  occur  even  at 
ordinary  temperatures.  Black  ferrous  sulphide  is  precipitated  from  ferrous 
solutions  by  alkaline  sulphides.  When  the  moist  sulphide  is  exposed  to 
the  air  it  oxidizes  to  ferrous  sulphate.  The  alkaline  sulphides  also  pre- 
cipitate ferrous  sulphide  from  ferric  salts,  but  the  latter  first  suffer  reduc- 
tion : 

2Fea3  -f  (NHd,S  = 2FeCl2  + 2NHW'l  -f  S, 

and 

FeCb  + (NHJgS  = FeS  + 2NH,C1. 


FERRIC  COMPOUNDS. 

Ferric  Sesquioxide  of  irofi^  exists  in  nature,  in  com- 

pact, massive  form,  as  red  hematite,  and  as  specular  iron,  in  dark-gray 
metallic  rhombohedra.  It  may  be  prepared  by  heating  the  iron  oxygen 
compounds  in  the  air,  and  is  obtained  on  a large  scale  by  the  ignition 
of  green  vitriol.  It  is  then  a dark-red  powder  {colcothar  or  caput 
7no7'tuu77i)  used  as  a paint  and  for  polishing  glass. 

Ferric  Hydroxide,  Fe(OH)3,  is  precipitated  by  alkalies  from  ferric 
solutions  as  a voluminous,  reddish-brown  mass.  On  boiling,  it  becomes 
more  compact,  gives  up  water,  and  is  converted  into  the  hydrate, 
Fe,^0(0H)^.  Many  iron  ores,  like  xanthosiderite,  Fe20(0H)4,  gdthite, 
Fe202(0H)2  (isomorphous  with  diaspore,  p.  350),  and  limonite, 
Fe^Og^OH)^,  are  analogous  compounds. 

Freshly  precipitated  ferric  hydroxide  is  soluble  in  a solution  of  ferric 
chloride  or  acetate.  When  such  a solution  is  subjected  to  dialysis,  the 
iron  salt  diffuses,  and  there  remains  a pure  aqueous  solution  of  ferric 
hydroxide.  All  of  the  latter  is  precipitated  as  a jelly  from  such  a solu- 
tion upon  the  addition  of  a little  alkali  or  acid. 

Ferrous-ferric  Oxide,  FCgO^  = Fe0.Fe20g,  occurs  in  nature  crys- 
tallized in  black  regular  octahedra — magnetite.  It  is  abundant  in 
Sweden  and  Norway,  and  in  the  Urals.  It  may  be  obtained  artificially 
by  conducting  steam  over  ignited  iron  (p.  397).  Magnetite  constitutes 
the  natural  loadstone. 


400 


INORGANIC  CHEMISTRY. 


Ferric  hydroxide,  like  otlier  sesciiiioxides,  is  a feelde  liase,  and  does 
not  yield  salts  with  weak  acids,  like  carbonic  or  siilpliurous  (p.  269). 

Ferric  salts  arise  by  the  solution  of  ferric  oxide  in  acids,  or  by  the 
oxidation  of  ferrous  salts  in  the  ])resence  of  free  acids  (best  by  chloric 
or  nitric  acids)  : 

2FeS(),  + 11,80,  + O = Fe,(S0,)3  d H/). 

Most  of  the  ferric  salts  have  a yellow-brown  color,  and  are  converted 
by  reduction  into  ferrous  salts  ; 

2FeCl3  t-  11,8  = 2FeCl2  -f  21101  -f  8. 

Ferric  Chloride,  FeCl,,.  It  is  obtained  in  aqueous  solution  by  con- 
ducting chlorine  into  a solution  of  ferrous  chloride: 

2FeCl2  -f  Cl,  2FeCl3. 

The  hydrate,  2FeCl3  -j-  3H2O,  remains  upon  evaporation.  It  is  a yel- 
low crystalline  mass,  readily  soluble  in  water,  alcohol,  and  ether.  When 
heated,  it  is  partly  decomposed,  hydrogen  escapes,  and  a mixture  of 
chloride  and  oxide  remains. 

Anhydrous  ferric  chloride  is  produced  by  heating  iron  in  a current  of 
chlorine  gas;  it  sublimes  in  brownish-green,  metallic,  shining,  six-sided 
prisms  and  scales,  which  deliquesce  in  the  air.  Ferric  chloride  boils  at 
280-285°.  Its  vai)or  density  between  320-440°  closely  approximates  the 
formula  Fe2Clg;  with  rising  temperature  it  diminishes  gradually,  and  from 
750-1050°  corresponds  to  the  simple  formula  FeCl,.  The  vapor,  however, 
])robably  does  not  consist  of  these  molecules,  because  at  440°  a decompo- 
sition into  ferrous  chloride,  FeCl,,  and  chlorine  commences  (Friedel  and 
Crafts,  Ber.  (1888)  21,  Ref.  580).  Recently  the  molecular  formula  FeCl, 
has  been  found  by  determining  the  rise  in  the  boiling  point  of  ethereal 
and  alcoholic  solutions  of  ferric  chloride. 

Ferric  chloride  solutions  can  take  up  large  quantities  of  ferric  hydrate. 
In  the  resulting,  dark-colored  solutions  ferric  hydroxychlorides  are 
present : 

nFeClg  -)-  mFe203  -j-  XH2O  ^Liquor  fcrri  oxychlorati). 

Ferric  Sulphate,  Fe2(SO,)3,  is  obtained  by  dissolving  the  oxide  in 
sulphuric  acid.  When  its  solution  is  evaporated,  it  remains  as  a white 
mass,  which  gradually  dissolves  in  water,  with  a reddish-brown  color.  It 
forms  alums  (p.  352)  with  alkaline  sulphates,  e.  g.  : 

Fe2(SO,)3.  K,80,  + 24H2O. 

Potassium  iron  alum. 

Ferric  Phosphate,  FePO,^  is  a yellowish-white  precipitate,  thrown 
out  of  ferric  solutions  by  sodium  phos])hate.  It  is  insoluble  in  water  and 
acetic  acid. 

Iron  Disulphide,  FeS,,  occurs  in  nature  as  iron  pyrites,  crystallized 
in  yellow,  metallic,  shining,  (^ctahedra  or  pentagonal  dodecahedra.  It  is 
employed  in  the  manufacture  of  sulphuric  acid  and  green  vitriol.  The 


CYANOGEN  DERIVATIVES  OF  IRON. 


401 


artificial  sulphide  can  be  prepared  in  many  ways.  It  leaves  ferrous  sul- 
phide when  it  is  strongly  heated  in  hydrogen. 


COMPOUNDS  OF  FERRIC  ACID. 

On  fusing  iron  filings  with  niter,  or  by  conducting  chlorine  into  potas- 
sium hydroxide,  in  which  ferric  hydroxide  is  suspended,  potassium  fer- 
rate, K2Fe04,  is  produced,  and  crystallizes  from  the  alkaline  solution  in 
dark-red  prisms.  This  salt  is  isomorphous  with  potassium  chromate  and 
sulphate.  It  dissolves  quite  easily  in  water ; but  the  dark-red  liquid  soon 
decomposes  with  separation  of  ferric  hydroxide  and  evolution  of  oxygen. 
The  free  acid  is  not  known,  as  it  immediately  breaks  down  when  liberated 
from  its  salts. 


CYANOGEN  DERIVATIVES  OF  IRON. 

When  potassium  cyanide  is  added  to  aqueous  solutions  of  the  ferrous  or 
ferric  salts,  the  cyanides,  Fe(CN)2  and  Fe(CN)3,  are  thrown  down  as  yel- 
lowish precipitates,  which  decompose  rapidly  in  the  air.  They  dissolve 
in  an  excess  of  potassium  cyanide  to  form  the  double  cyanides,  Fe(CN)2.- 
4KCN  and  Fe(CN)3. 3KCN.  When  acids  are  added  to  strong  solutions 
of  these  salts  the  hydrogen  compounds,  H4Fe(CN)g,  and  H3Fe(CN)g  sepa- 
rate. Like  the  halogen  hydrides  they  are  acids  and  form  salts  by  exchang- 
ing their  hydrogen  for  metals.  The  iron  and  the  cyanogen  group  in  these 
salts  and  in  the  free  acids  cannot  be  detected  by  the  usual  reagents  {e.  g., 
the  iron  is  not  precipitated  by  the  alkalies).  It  is  supposed  that  com- 
pound groups  of  peculiar  structure  are  present  in  these  double  cyanides, 
and  that  they  conduct  themselves  like  the  halogens.  The  group,  FeCyg, 
in  the  i^rrous  compounds  is  called  ferrocyanogen,  that  of  FeCyg  in  the 
ferr/V,  ferricyanogen.  (Nothing  is  known  in  regard  to  the  molecular 
magnitude  of  these  compounds;  Cy  = CN).  The  ferro-  behave  toward 
the  ferri-compounds  the  same  as  the  ferrous  toward  the  ferric  salts ; oxi- 
dizing agents  convert  the  former  into  the  latter,  and  reducing  agents 
transform  the  latter  into  the  former : 


and 


K,Fe(CN)g  + Cl  = K3Fe(CN)g  + KCl 
K3Fe(CN)g  -f  KOH  + H = K4Fe(CN)e  -f  H^O. 


Cobalt,  manganese,  chromium  and  the  platinum  metals  afford  similar 
cyanides  (pp.  377,  407). 

Metallic  acid  radicals  such  as  are  known  to  exist  in  the  chromates,  manganates,  per- 
manganates and  ferrates  are  assumed  to  be  present  in  the  ferro-  and  ferricyanides.  Such 
radicals,  with  a basic  character,  have  been  encountered:  the  radicals  (WgCh)!!  and 
(MogCb)!!  which  j)lay  the  role  of  a bivalent  metal  towards  the  halogens  in  the  com- 
pounds— chlor-tungsten  hydroxide  and  chlor-molybdenum  hydroxide  (p.  382). 

All  of  these  radicals  have  this  in  common — the  metal  present  in  the  radical  is  no 
34 


402 


in()R(;anic  chemistry. 


longer  (letccled  l)y  the  reagents  to  wliieli  it  responds  wlien  existing  as  metal  in  its  salts, 
'This  it  will  oidy  do  after  the  radical  has  been  destroyed.  If  the  radieals  eonlain  halogens 
they  too  will  behave  like  the  metals.  'This  failure  to  resi)ond  to  the  usual  reagents  is  an 
indieation  that  the  metal  or  halogen  belongs  to  a compound  radical  and  docs  ncjt  act  ns  an 
independent  part  in  the  molecule.  As  mentioned  on  p.  270,  salts,  bases  and  acids  are 
regarded  as  compounds  of  two  members  when  they  are  resolved  in  acjueous  solutiem  into 
two  kinds  of  ions.  Numerous  transpositions  and  particularly  those  of  analytical  value 
occur  between  ions,  hence  they  are  briefly  termed  “ion  reactions.”  I'.very  ion,  be  it 
simple  or  compound,  is  characterized  by  reactions  which  belong  to  it  alone  and  by  which 
it  is  detected.  Hence  in  general  we  cannot  speak  of  the  reactions  of  the  elements,  e.  jif., 
of  iron,  when  search  is  being  made  for  its  presence,  Init  we  must  rather  consider  the  con- 
dition, the  combination  form,  the  nature  of  tlie  ion,  for  whicli  the  reaction  has  value. 
There  are  no  general  reagents  or  reactions  for  iron,  but  only  those  for  metallic  iron,  or 
for  ferrous,  ferric,  and  ferrocyanide  comi)ounds,  etc.  ; no  general  reaction  for  chlorine, 
but  only  such  as  answer  for  free  chlorine,  hydrogen  chloride,  cldoric  acid,  percldoric 
acid,  etc. — reactions  of  the  ions  according  to  the  theory  of  electrolytic  dissociation,  d'his 
is  expressed  in  the  following  formulas  which  are  divided  by  lines  into  their  ions  : 

FeiCk,  K^KFeCyg)  II  Cl  HKClOj)  Mnl(SO,)  K^^MnOJ 
MoiCl,  (Mo.p,)|Cl2  Cr2!(SOj3  K^KCrOJ. 

Potassium  Ferrocyanide,  Yellow  prussiate  of  potash,  K^Fe(CN)g, 
is  produced  by  the  action  of  potassium  cyanide  upon  iron  compounds,  or 
upon  free  iron  (in  which  case  the  oxygen  of  the  air  or  water  takes  part). 
It  is  prepared  commercially  by  igniting  carbonized  nitrogenous  animal 
matter  (blood,  horns,  hoofs,  leather  offal,  etc.)  with  potashes  and  iron. 

In  this  operation,  the  carbon  and  nitrogen  of  the  organic  matter  com- 
bine with  the  potassium  of  the  potashes  to  form  potassium  cyanide,  while 
the  sulphur  present  forms  iron  sulphide  with  the  iron.  (By  means  of 
alcohol,  potassium  cyanide  can  be  extracted  from  the  fusion.)  Upon 
treating  the  fusion  with  water,  the  potassium  cyanide  and  iron  sulphide 
react  upon  one  another,  and  potassium  ferrocyanide  results  and  is  puri- 
fied by  crystallization  : 

FeS  + 6KCN  = K,Fe(CN)6  + K^S. 

In  Germany,  at  present,  yellow  prussiate  of  potash  is  manufactured 
exclusively  from  the  material  used  in  purifying  gas  (p.  105),  which  con- 
tains the  greater  part  of  the  cyanogen  of  the  crude  gas  in  the  form  of 
Prussian  blue  and  ammonium  sulphocyanide.  The  first  is  transformed 
by  lime  into  ferric  oxide  and  calcium  ferrocyanide,  which  in  aqueous 
solution  is  changed  by  potassium  chloride  to  calcium  chloride  and  very 
sparingly  soluble  potassium  calcium  ferrocyanide  : 

Ca,Fe(CN)e  + 2KCI  =3  CaK.,Fe(CN)6  + CaCh. 

Potashes  are  finally  used  to  convert  the  insoluble  salt  into  calcium  carbon- 
ate and  potassium  ferrocyanide. 

It  crystallizes  from  water  in  large,  yellow,  monoclinic  prisms,  having 
three  molecules  of  water,  and  soluble  in  3-4  parts  of  water.  The  crystals 
lose  all  their  water  at  100°,  and  are  converted  into  a white  })owder.  At  a 
red  heat  the  ferrocyanide  breaks  down  into  potassium  cyanide,  nitrogen, 
and  iron  carbide  (FcG^).  When  the  salt  is  warmed  with  dilute  sulphuric 


CYANOGEN  DERIVATIVES  OF  IRON. 


403 


acid,  half  of  the  cyanogen  escapes  as  hydrogen  cyanide  ; concentrated 
sulphuric  acid  decomposes  it,  according  to  the  following  equation  : 

K,Fe(CN)6  + 6H2SO,  + 6H,0  = FeSO,  + aK^SO^  + 3(NH4),SO,  + 6CO. 

When  strong  hydrochloric  acid  is  added  to  a concentrated  solution  of 
potassium  ferrocyanide  hydrogefi  ferrocyanide,  H^FeCyg,  separates  as  a 
white  crystalline  powder,  which  soon  turns  blue  in  the  air.  It  is  an  acid. 
Its  salts  with  the  alkali  and  alkaline  earth  metals  are  very  soluble  in  water. 
The  sodium  salt  crystallizes  with  difficulty.  The  salts  of  the  heavy  metals 
are  insoluble  in  water,  and  are  obtained  by  double  decomposition.  When 
potassium  ferrocyanide  is  added  to  the  solution  of  a ferric  salt  a dark- 
blue  cyanide 

Fe^(Cy)ig  = (Fe2)2(FeCyg)3, 

called  Prussian  blue,  is  precipitated  : 

3K,FeCyg  + 4Fe2Clg  = (Fe2)2(FeCyg)3  + 12KCI. 

This  color  was  discovered  accidentally  by  Diesbach,  of  Berlin,  in  1704. 

It  is  the  ferric  salt  of  hydroferrocyanic  acid;  and  if  potassium  or 
sodium  hydroxide  is  poured  over  it,  it  is  converted  into  ferrocyanide  of 
potassium  and  ferric  hydroxide; 

VI  VI 

(Fe2)2(FeCyg)3  + 12KOH  = 3K^FeCyg  -f-  4Fe(OH)3. 

Potassium  ferrocyanide  produces  a reddish-brown  precipitate  of  cop- 
per ferrocyanide,  CiqFeCyg,  in  copper  solutions ; and  in  ferrous  solutions  a 
white  precipitate,  which  on  exposure  to  the  air  rapidly  becomes  blue  in 
color. 


Oxidizing  agents  convert  the  ferro-  into  potassium  ferricyanide, 
K^FeCyg,  red  prussiate  of  potash.  This  conversion  is  most  conveniently 
effected  by  conducting  chlorine  into  the  solution  of  the  yellow  prussiate  : 

aK^FeCyg  -j-  Clg  = 2K3FeCyg  -j-  2KCI. 

The  quadrivalent  ferrocyanogen  group,  FeCyg,  is  then  changed  to  the 
trivalent  ferricyanogen  group,  FeCyg  (p.  401).  Cy  = CN. 

The  red  prussiate  crystallizes  from  water  in  red  rhombic  prisms.  The 
free  hydroferricyanic  acid,  H3Fe(CN)g,  is  precipitated  upon  the  addition 
of  concentrated  hydrochloric  acid.  It  is  rather  unstable. 

With  ferrous  solutions  potassium  ferricyanide  yields  a dark-blue  pre- 
cipitate, FegCyij  = FegFcaCyii,  very  similar  to  Prussian  blue,  and  called 
Turnbuir  s blue : 

2K3FeCyg  -f  3FeSO^  = FegCy,^  + 3^280,. 

So  far  as  its  formation  is  concerned,  'rurnbull’s  blue  is  the  ferrous  salt  of  hydroferri- 
III  II 

cyanic  acid,  [Fe(CN)g] Fcg.  Although  liot  alkalies  decompose  it  into  potassium  ferro- 
cyanide and  ferrous-ferric  hydroxide,  this  is  most  probably  due  to  the  fact  that  the  potassium 


404 


INORGANIC  CIIKMIS'I’RY. 


ferricyanide  and  ferrous  liydroxidc,  which  arc  i)roduccd  at  first,  arc  traiisjjoscd  by  the 
alkali  into  potassium  ferrocyanide  and  fcrrous-fcrric  hydrate  : 

2Fe(CN),K3  -I-  2KOII  4 3l''e{OII),  ^ 2K,Fc(CN),  + 2Fc(OII)3  t Fc(()II),. 

Very  probably  in  the  i)recipitation  potassium  ferricyanide  and  the  ferrous  salt  transpose 
into  potassium  ferrocyanide  and  a ferrous-ferric  salt  ; then  'I'urnbuirs  blue  would  have  to 

III  II  IV 

be  regarded  as  the  ferric-ferrous  salt  of  hydroferrocyanic  acid — J‘'e.4‘'e[Fe((:N),.].^.  'I'urn- 
bull’s  blue  is  the  principal  constituent  of  commercial  J’russian  blue,  which  also  contains, 
because  of  its  preparation  from  impure  materials,  ferriferrocyanide,  Fe7(CN),y.  '1  he 
latter  largely  comprises  Parisian  blue. 

Potassium  ferricyanide  does  not  cause  precipitation  in  ferric  solutions, 
pjy  these  reactions,  ferric  salts  may  be  readily  distinguished  from  the 
ferrous.  Potassium  sulphocyanide  (KCNS)  produces  a dark- red  colora- 
tion in  ferric  solutions,  while  it  leaves  the  ferrous  unaltered. 

Iron,  like  nickel  (]).  233),  combines  with  carbon  monoxide  to  a vola- 
tile, gaseous  body — Iron  Carbonyl  [Mond  and  Quincke,  Per.  (1891) 
24,  2248;  Perthelot,  Compt.  rend.  (1891)  112,  1343],  which,  from  its 
composition,  is  probably  Iron  Tetracarbonyl,  Fe(CO)^.  This  substance 
is  produced  by  conducting  carbon  monoxide  over  very  finely  divided 
iron  at  40-80°,  or  under  a pressure  of  eight  atmospheres  [Roscoe  and 
Scudder,  Per.  (1891)  24,  3843].  It  is  decomposed  with  the  formation 
of  an  iron  mirror  when  it  is  passed  through  a glass  tube  heated  to  200- 

350°- 

Iron  pentacarbonyl,  Fe(C0)5,  hepiacarbonyl,  Fe.2(CO)y,  have  been  prepared. 

The  first  is  a yellow  liquid,  boiling  at  103°,  from  which  the  second  separates,  on  expos- 
ure to  the  light,  in  golden  crystals.  Carbon  monoxide  is  evolved  at  the  same  time. 


2.  NICKEL. 

Ni  = 58.7.* 

Nickel  exists  in  native  condition  in  meteorites;  its  most  important 
ores  are  niccolite,  NiAs,  and  gersdorffite,  NiSa . NiAsj  (constituted  like 
cobaltite).  The  arsenical  ores  are  now  of  little  importance  in  the  nickel 
industry.  The  chief  sources  of  the  metal  are  at  present  nickeliferous 
])hyrrotite  and  the  nickel  silicates  (Canada,  Norway  and  New  Caledonia). 
Clarnierite — a New  Caledonian  mineral  — a silicate,  containing  also  iron, 
calcium  and  magnesium,  may  be  especially  noticed  in  this  connection. 
Nickel  is  almost  always  accompanied  in  its  ores  by  cobalt,  and  vice  versa, 
cobalt  usually  by  nickel.  The  isolation  of  the  latter  from  its  ores  and 
from  speiss-cobalt  (p.  406)  is  very  complicated.  Nickel  usually  appears 
in  cfimmerce  in  cubical  forms,  which  in  addition  to  the  chief  ingredient 
alvvays  contain  some  copi)er,  bismuth,  and  other  metals.  Chemically 
))ure  nickel  is  procured  by  igniting  the  oxalate  in  a current  of  hydrogen. 
Nickel  is  almost  silver-white  in  color  and  is  very  lustrous,  and  very 


* Sc(; 'I'll.  \V.  Richards  and  A.  S.  ('ushman,  Z.  f.  anorg.  Chem.  20  (1899),  352. 


NICKEL. 


405 


tenacious.  Its  specific  gravity  varies  from  8.8  to  9.1.  It  fuses  at  a some- 
what lower  temperature  than  iron,  and  like  it  is  attracted  by  the  magnet. 
It  is  not  altered  in  the  air;  it  dissolves  with  difficulty  in  hydrochloric  and 
sulphuric  acids,  but  readily  in  nitric  acid. 

Its  derivatives  are  almost  exclusively  of  the  ous  form,  NiX.^;  nickelic 
oxide  behaves  like  a peroxide,  and  does  not  form  corresponding  salts. 

Nickelous  Hydroxide,  Ni(OH).^,  is  a bright  green  precipitate  pro- 
duced by  alkalies  in  nickelous  solutions.  It  dissolves  in  ammonium 
hydroxide,  with  a blue  color.  When  heated  it  passes  into  gray  nickelous 
oxide,  NiO. 

Nickelous  Chloride,  NiCl^ -f  bH.^O,  consists  of  green,  monoclinic 
prisms.  When  heated  they  lose  water  and  become  yellow. 

Nickelous  Cyanide,  Ni(CN)2,  is  preci})itated  by  potassium  cyanide 
as  a green-colored  mass  from  nickel  solutions.  It  is  soluble  in  excess  of 
the  precipitant.  The  double  cyanide,  NiCyj. 2KCy -|- H2O,  crystallizes 
from  the  solution.  This  salt  is  readily  decomposed  by  acids.  Cyanogen 
compounds  of  nickel,  constituted  like  those  of  iron  and  cobalt,  are  not 
known. 

Nickelous  Sulphate,  NiSO^ -j- 7H2O,  ■ appears  in  green,  rhombic 
prisms,  isomorphous  with  the  sulphates  of  the  magnesium  group,  and 
forms  analogous  double  salts. 

Nickelous  Sulphide,  NiS,  is  precipitated,  black  in  color,  by  alkaline 
sulphides  from  nickel  solutions. 

Nickelic  Oxide,  Ni203  and  Hydroxide,  Ni2(OH)g,  are  perfectly 
similar  to  the  corresponding  cobalt  salts ; when  warmed  with  hydrochloric 
acid  they  liberate  chlorine. 

Nickel  Tetracarbonyl,  Ni(CO)4,  see  pp.  234,  404. 

Nickel  is  used  for  certain  alloys  Argentan  consists,  ordinarily,  of  50 
per  cent,  of  copper,  25  per  cent,  of  nickel  and  25  per  cent,  of  zinc.  The 
white  color,  the  hardness  and  the  power  of  receiving  a higher  polish  in- 
crease with  the  increase  in  the  nickel  content.  The  iron-nickel  alloys 
have  great  technical  importance  (nickel  steel).  The  German  nickel 
coins  in  which  we  have  evidence  of  the  great  coloring  power  of  nickel 
consist  of  75  per  cent,  of  copper  and  25  per  cent,  of  nickel.  Nickel 
alloys  are  used  to  make  electrical  resistances;  their  conductivity  is  slight 
and  little  depends  on  this  at  the  ordinary  temperature  {niaiiganin : 84 
per  cent,  of  copper,  4 per  cent,  of  nickel,  12  per  cent,  of  manganese; 
consfaiiian : 60  per  cent,  of  copper,  40  per  cent,  of  nickel,  etc.).  At 
present,  cast-iron  ware  is  coated  with  a layer  of  nickel  to  prevent  it  from 
rusting  and  to  imj)art  to  it  a beautiful  white  surface.  This  is  accomplished 
by  electroplating,  or  by  boiling  the  iron  ware  in  a solution  of  zinc 
chloride  and  nickel  sulphate. 

In  the  electrolytic  method  a solution  of  the  double  sulphate  of  nickel 
and  ammonium  is  emjdoyed  ; the  positive  electrode  consists  of  a pure 
nickel  plate,  while  the  object  to  be  coated  forms  the  negative  electrode. 


4o6 


INORGANIC  CIIEMISIRY. 


3.  COBALT. 

Co  = 59  * 

Cobalt  occurs  in  nature  as  sinallile  (CoAs^)  and  cobaltite  (CoAs.^.  CoSj). 
The  metal  is  obtained  by  the  ignition  of  cobaltous  oxide  with  carbon,  or 
in  a current  of  hydrogen.  It  has  a reddish-white  color  and  strong  luster, 
is  very  tenacious,  and  fuses  with  difficulty.  Its  specific  gravity  is  8.5.  It 
is  attracted  by  magnets,  but  to  a less  degree  than  iron.  It  is  not  altered 
by  the  air  or  water.  It  is  only  slightly  attacked  by  hydrochloric  and 
sulphuric  acids  ; nitric  acid  dissolves  it  readily,  forming  cobaltous  nitrate. 

The  predominating  compounds  have  the  form  CoX.^,  and  are  called 
cobaltous.  They  are  very  stable,  and  generally  isomorphous  with  the 
corresponding  ferrous  salts,  'bhe  hydrous  cobaltous  compounds  have  a 
reddish  color,  the  anhydrous  are  blue. 


COBALTOUS  COMPOUNDS. 

Cobaltous  Chloride,  C0CI2,  is  obtained  by  the  solution  of  cobaltous 
oxide  in  hydrochloric  acid,  and  crystallizes  with  six  molecules  of  water 
in  red  monoclinic  prisms.  When  heated,  it  loses  water,  and  becomes 
anhydrous  and  blue  in  color.  Characters  made  with  this  solution  upon 
paper  are  almost  invisible,  but  when  warmed  they  become  distinct  and 
blue  (sympathetic  ink). 

Cobaltous  Hydroxide,  Co(OH)2,  is  a reddish  precipitate  produced 
by  the  alkalies  in  hot,  cobaltous  solutions.  When  exposed  to  the  air,  it 
is  colored  brown  by  oxidation.  Basic  salts  are  precipitated  from  cold 
solutions;  these  dissolve  with  a blue  color  in  an  excess  of  concentrated 
alkalies.  When  heated  out  of  air  contact,  the  hydroxide  passes  into 
green  cobaltous  oxide,  CoO. 

Cobaltous  Sulphate,  CoSO^  -j-  yH^O,  crystallizes  in  dark-red 
monoclinic  prisms;  the  hydrated  suljffiate,  CoSO^  -j-  6H2O,  separates 
from  hot  solutions.  It  is  isomorphous  with  ferrous  sulphate,  and  yields 
double  salts  with  alkaline  sulphates  (p.  398). 

Cobaltous  Nitrate,  Co(N03)2  + 6H2O,  forms  red  deliquescent 
prisms. 

Cobaltous  Sulphide,  CoS,  is  a black  precipitate,  produced  in  neu- 
tral cobalt  solutions  by  alkaline  sulphides.  It  is  insoluble  in  dilute 
acids. 

Cobalt  Silicates. — When  glass  is  fused  with  a cobalt  compound  it  is 
colored  a dark  blue,  and  when  reduced  to  a powder  is  used  as  a pigment 
under  the  name  of  smalt. 

Smalt  is  j)repared  comitiercially  by  fusing  cobalt  ores  with  potashes  and  quartz.  The 
cobalt  forms  a silicate  (smalt)  with  the  silica  and  potassium,  while  the  other  metals,  such 
as  bismuth,  arsenic,  and  csi^ecially  nickel,  accompanying  it  in  its  ores,  are  thrown  out  as 
a metallic  mass.  'I'his  is  called  si)eiss-cobalt  and  serves  for  the  preparation  of  nickel. 


* See  (d.  Winkler,  Z.  f.  anorg.  Ch.  17  (1S98),  236. 


COBALTIC  COMPOUNDS. 


407 

On  igniting  cobalt  oxide,  CO2O3,  with  alumina,  a dark-blue  mass  is 
produced — cobalt  uli7-a7tia7-ine  or  The7ia7'd^ s blue.  When  zinc  oxide  and 
cobalt  oxide  are  ignited  a green  color — greeTt  cmTtabar  or  Ri7i7naft7i' s 
gree7i — is  obtained. 


COBALTIC  COMPOUNDS. 

Cobaltic  Oxide,  CO2O3,  is  left  as  a black  powder  on  the  ignition  of 
cobaltous  nitrate.  It  becomes  cobaltous-cobaltic  oxide,  COgO^,  at  a red 
heat,  and  cobaltous  oxide  at  a white  heat.  The  hydroxide,  Co2(OH)j., 
separates  as  a dark-brown  powder,  if  chlorine  be  passed  through  an  alka- 
line solution  containing  a cobaltous  salt. 

A cobaltous  salt  is  produced  and  oxygen  set  free,  when  sulphuric  acid 
acts  upon  the  oxide  or  the  hydroxide.  Chlorine  is  generated  when  it  is 
heated  with  hydrochloric  acid  : 

C02O3  -f  6HC1  = 2C0CI2  + 3H2O  + CI2. 

The  cobaltic  hydroxide  dissolves  in  dilute,  cold  hydrochloric  acid, 
with  scarcely  any  liberation  of  chlorine  ; the  solution  probably  contains 
cobaltic  chloride,  C0CI3,  which  decomposes  into  cobaltous  chloride  and 
chlorine  on  evaporation. 

Cobaltous-cobaltic  Oxide,  0030^=00203.000,  corresponding 
to  magnetite,  FegO^,  is  formed  upon  the  ignition  of  the  oxygen  cobalt 
derivatives  in  the  air,  and  is  a black  powder. 


It  is  noteworthy  that  cobalt  is  capable  of  yielding  complex  derivatives 
in  which  it  appears  to  be  in  union  with  the  groups  NO2,  ON  or  NH3, 
forming  peculiar  radicals.  One  of  the  most  interesting  of  these  is 
potassium  cobaltic  nitrite  (Fischer’s  salt). 

When  potassium  nitrite,  KNOg,  is  added  to  a cobaltous  solution  acidi- 
fied with  acetic  acid,  nitric  oxide  is  set  free,  and  in  course  of  time 

Co(N02)3.3KN02+  nH20, 

a double  salt,  separates  as  a yellow  crystalline  powder.  This  should  be 
viewed  as  the  potassium  salt  of  a hydronitroso-cobaltic  acid  : 

H3Co(N02V 

This  reaction  is  very  characteristic  for  cobalt,  and  serves  to  separate  it 
from  nickel. 

Ammonia-cobalt  Compounds. — Numerous  cobaltamines  are  known.  On  adding 
ammonium  hydroxide  to  a cobaltous  chloride  solution,  the  precipitate  first  formed  dissolves 
in  the  excess  of  the  reagent,  and  when  this  liquid  is  permitted  to  stand  exposed  to  the  air, 
the  color,  which  is  brown  at  first,  gradually  passes  into  red.  On  adding  concentrated 
hydrochloric  acid  to  this  solution,  a brick-red,  crystalline  powder,  of  the  composition 

C02CI3.10NH3  -f  2II2O, 


4o8 


INORGANIC  CHEMISTRY, 


called  roseocobaltic  chloride^  is  precipitated.  If,  however,  the  red  solution  be  boiled  with 
hydrochloric  acid,  a red  j)owder,  piirpureocobaltic  chloride^ 

COjClg.  lONIIg, 

separates  out.  If  the  ammoniacal  red  solution  contain  much  ammonium  chloride,  hydro- 
chloric acid  will  precipitate  a yellowish-brown  compound — liUeocobaltic  chloride, 

Co2Cl6.’i2XIl3. 

These  derivatives  are  supposed  to  contain  compound  basic  radicals  in  which  cobalt,  halo- 
gens and  ammonia  are  all  })resent. 

'I'he  other  salts  of  cobalt,  such  as  the  sulphate  and  nitrate,  yield  similar  compounds,  : 

Co3(N03)e.ioNIl3, 

roseocobaltic  nitrate. 

Cyanogen  Cobalt  Compounds. — In  solutions  of  cobaltous  salts,  potassium  cyanide 
produces  a bright  brown  precipitate  of  cobaltocyanide, 

Co(CN)3, 

soluble  in  an  excess  of  the  reagent.  The  solution  absorbs  oxygen  from  the  air,  and  is 
converted  into  potassium  cobalticyanide, 

K3Co(CN)„ 

corresponding  to  potassium  ferricyanide.  When  the  solution  is  evaporated  the  cobalti- 
cyanide  crystallizes  in  colorless  rhombic  prisms,  very  soluble  in  water.  Sulphuric  acid 
precipitates  hydrogen  cobalticyanide, 

Il3Co(CN)„ 

from  the  concentrated  solution.  This  acid  crystallizes  in  needles. 


GROUP  OF  THE  PLATINUM  METALS. 

Besides  platinum,  this  group  comprises  palladium,  rhodium,  ruthenium, 
osmium,  and  iridium — the  constant  companions  of  the  first  in  its  ores. 
On  p.  392  we  observed  that  these  metals  are  divided  into  two  groups: 
{hegrou/>  of  light  platmiim  tnetals,  and  group  of  heavy  platinum  metals. 
The  latter  have  higher  atomic  weights  and  specific  gravities: 

Ru,  101.7  Rh,  103.0  Pd,  106  Os,  191  Ir,  193.0  Pt,  194.8 

Sp.gr.  “ 12.3  “ 12. 1 “ II. 5 “ 22.48  “ 22.4  “ 21.5 

At.  vol.  “ 8.3  “ 8.6  “ 9.1  “ 8.4  “ 8.6  “ 9.0. 

The  relations  of  the  metals  of  these  two  groups  to  each  other  are  per- 

fectly similar  to  those  of  the  iron  group.  Osmium  and  ruthenium,  like 
iron,  have  a gray  color,  fuse  with  difficulty,  and  are  readily  oxidized  in 
the  air.  Palladium  and  ])latinum,  on  the  other  hand,  have  an  almost 
silver-white  color  like  nickel,  are  more  easily  fusible,  and  are  not  oxidized 
by  oxygen.  In  chemical  respects  osmium  and  ruthenium,  like  iron,  also 
show  a metalloidal  nature,  inasmuch  as  their  highest  oxygen  compounds 


GROUP  OF  THE  PLATINUM  METALS.  409 

form  acids.  Their  derivatives  show  a complete  parallelism  with  those  of 
iron  : 


II 

Ill 

IV 

VI 

OsO 

OSjOg 

OSO2 

(OSO3) 

Osmous 

Osmic 

Osmium 

Osmium 

oxide. 

oxide. 

dioxide. 

trioxide. 

RuO 

RUjOj 

Ruthenic 

RUO2 

(RUO3) 

Ruthenious 

Ruthenium 

Ruthenium 

oxide. 

oxide. 

dioxide. 

trioxide. 

The  acid  oxides  OsOg  and  RuO,  are  unknown,  but  the  corresponding 
acids,  H20sO^  (osmic  acid)  and  H2RuO^  (ruthenic  acid),  and  their  salts 
have  been  obtained.  Besides  the  derivatives  already  mentioned  we  find 
that  osmium  and  ruthenium  are  capable  of  still  higher  oxidation,  yielding 
OsO^,  per-osmic  anhydride,  and  RuO^,  per-ruthenic  anhydride — which  is 
not  the  case  with  iron  ; in  these  compounds  the  metals  appear  to  be  octads, 
yet  these  oxides  do  not  form  corresponding  acids  or  salts. 

Rhodium  and  iridium,  like  cobalt,  do  not  yield  acid-like  derivatives. 
Their  salts  correspond  to  the  forms : 

II  III  IV 

RhO  Rh203  Rh02 

Rhodous  Rhodic  Rhodium 

oxide.  oxide.  dioxide. 

The  rhodic  compounds  are  the  more  stable  derivatives. 

Palladium  and  platinum,  finally,  are  relatively  of  more  basic  nature,  as 
their  ous  derivatives,  PdX2  and  PtX2,  are  proportionally  more  stable  than 
the  ic  forms,  PdX^  and  PtX^.  Palladium  also  forms  a lower  oxide,  palla- 
dium suboxide,  Pd20,  in  which  it  approaches  silver. 


The  platinum  metals  are  found  in  nature  almost  exclusively  in  the  so- 
called  platinum  ore,  which  usually  occurs  in  small  metallic  grains  in  accu- 
mulated sands  of  a few  regions  (in  California,  Australia,  the  island  of 
Sumatra,  and  especially  in  the  Urals).  The  platinum  ore,  like  that  of 
gold,  is  obtained  by  the  elutriation  of  the  platiniferous  sand  with  water, 
whereby  the  lighter  particles  are  carried  away.  Platinum  ore  usually  con- 
tains 50-80  percent,  of  platinum,  besides  palladium  (to  2 per  cent.),  irid- 
ium (to  7 per  cent.),  osmium  (i^  per  cent.),  and  ruthenium  (i)4  pei' 
cent.),  and  different  other  metals,  as  gold,  copper,  and  iron. 

The  separation  of  the  platinum  metals  is  generally  executed  in  the  fol- 
lowing manner:  The  gold  is  first  removed  by  dilute  aqua  regia.  Then 
the  ore  is  treated  with  concentrated  aqua  regia,  when  platinum,  palladium, 
rhodium,  ruthenium,  and  a portion  of  iridium  are  dissolved.  Metallic 
grains  or  leaflets,  an  alloy  of  osmium  and  iridium — platinum  residues — 
remain.  Ammonium  chloride  is  then  added  to  the  solution  and  platinum 
and  iridium  precipitated  as  double  salts.  When  the  precipitate  is  ignited, 
a spongy  mass  of  iridium  bearing  platinum  (platinum  sponge)  is  obtained, 
which  is  applied  directly  in  the  manufacture  of  platinum  vessels.  The 
filtered  solution  from  the  insoluble  chlorides  contains  palladium,  rhodium, 
35 


410 


INORGANIC  CHEMISTRY. 


and  ruthenium,  which  are  thrown  down  as  a metallic  powder  by  iron; 
their  further  sejiaration  is  then  effected  in  various  ways. 

Formerly  spongy  })latinum  was  employed  almost  exclusively  for  the 
manufacture  of  })latinum  objects  ; itwas  pressed  into  moulds,  then  ignited 
and  hammered  out.  Now  the  fusibility  of  jdatinum  in  the  oxyhydrogen 
flame  is  resorted  to,  and  the  fused  metal  run  into  moulds. 

Platinum  containing  both  iridium  and  rhodium  may  be  fused  directly 
out  of  the  platinum  ore  by  means  of  the  oxyhydrogen  blowpi])e.  The 
greater  ])ortion  of  the  osmium  and  ruthenium  is  consumed  in  this  opera- 
tion. The  jiresence  of  iridium  and  rhodium  makes  platinum  harder  and 
less  readily  attacked  by  many  reagents. 


RUTHENIUM  AND  OSMIUM. 

Ru  101.7.  Os  = 191. 

Ruthenium,  discovered  by  Claus  in  1845,  has  a steel-gray  color  ; it  is  very  hard,  brittle, 
and  difficultly  fusible  (at  about  1800°).  When  pulverized  and  ignited  in  the  air  it  oxidizes 
to  RuO,  Ru^03  and  RuO.^.  It  is  insoluble  in  acids,  and  only  slowly  dissolved  by  aqua 
regia.  When  fused  with  potassium  hydroxide  and  nitrate,  it  forms  potassium  ruthenate, 
K2l^^i04- 

Ruthenium  heated  in  chlorine  gas  yields  ruthenium  dichloride^  RuClj,  a black  pow- 
der, insoluble  in  acids.  The  trichloride,  RuCl,,  is  obtained  by  the  solution  of  Ru2(0ing 
in  hydrochloric  acid,  and  is  a yellow,  crystalline  mass,  which  deliquesces  in  the  air.  It 
yields  crystalline  double  chlorides  with  potassium  and  ammonium  chlorides,  e.  g., 
RuClg.KCl.  The  tetrachloride,  RuCh,  is  only  known  in  double  salts.  Ruthenious 
oxide,  RuO,  the  sesquioxide,  RU2O3,  and  dioxide,  RUO2,  are  black  powders,  insoluble  in 
acids,  and  are  obtained  when  ruthenium  is  roasted  in  the  air. 

The  hydroxides,  RujlOH)^  and  Ru(0H)4,  are  produced  by  the  action  of  the  alkalies 
upon  the  corresponding  chlorides,  and  are  very  readily  soluble  in  acids.  Ruthenic  acid, 
H2RUO4,  is  not  known  in  a free  condition.  Its  potassium  salt,  K2RuO^,  is  formed  by 
fusing  the  metal  with  potassium  hydroxide  and  niter  or  by  digesting  the  tetroxide  with 
dilute  alkali  at  60°  until  the  evolution  of  oxygen  ceases : 

RuO^  -f-  2KOH  = KgRuO^  HgO  O. 

It  crystallizes  in  black  prisms  with  a green  luster.  It  contains  one  molecule  of  water  of 
crystallization.  It  absorbs  moisture  and  carbon  dioxide  with  avidity  from  the  air.  Its 
deep  orange-red,  dilute,  aqueous  .solution  gradually,  on  exposure  to  the  air,  becomes 
green  and  a perruthenate  results.  At  the  same  time  a black  oxide  of  the  formula  RU2O5 
separates.  This  decomposition  is  more  rapid  when  the  solution  is  acted  upon  with  car- 
bon dioxide,  dilute  acids,  chlorine  or  bromine.  Potassium  ruthenate  therefore  behaves 
like  potassium  manganate  (p.  390).  When  chlorine  is  conducted  through  a concentrated 
.solution  of  ])otassium  ruthenate,  or  when  ruthenium  is  roasted  at  1000°  in  a current  of 
oxygen,  ruthenium  tetroxide,  RuO^,  and  dioxide  are  formed.  By  the  chlorine  method 
it  volatilizes  and  may  be  collected  in  a well-cooled  receiver.  It  consists  of  a gold-yellow 
crystalline  mass.  It  fu.ses  at  25.5°,  therefore  in  the  hand,  to  a deep  orange-red  colored 
li(|uid.  It  readily  sublimes  in  large,  yellow  transparent  ciystals  with  an  orange-yellow 
reflex.  At  the  ordinary  pressure  it  does  not  boil  at  106°,  but  it  decomposes  with  ex- 
plosion at  108°.  It  can  be  gasified  without  decomposition  under  diminished  pressure. 
Its  vapors  then  correspond  to  the  formula  RuO^.  'bhe  tetroxide  dissolves  slowly  in 
water  with  a golden-yellow  color,  but  does  not  yield  a hydrate.  Black  compounds  of 
varying  comi)()sition  se[)arate  from  this  solution.  Potassium  perruthenate,  KRuO^,  is 
furni'al  along  with  the  more  iiolublc  julhcnalc  on  adding  the  tdtroxidc  to  caustic  potash. 


RHODIUM  AND  IRIDIUM. 


41I 

It  crystallizes  in  black  octahedra  with  metallic  luster;  they  dissolve  in  water  with  a deep- 
green  color  [see  Debray  and  Joly,  Jahresber.  der  Chemie  (1888),  669,  672]. 

Osmium  is  very  much  like  ruthenium.  It  is  not  even  fusible  in  the  oxyhydrogen 
flame  ; it  only  sinters  together.  According  to  Violle  it  fuses  at  2500°.  Of  all  the  plati- 
num metals  it  is  the  most  difficult  to  fuse,  but  is  the  most  easily  oxidized.  Covipact 
osinhivi  is  the  heaviest  metal  specifically  (specific  gravity  22.48)  and  is  insoluble  in  aqua 
regia.  Reduced  to  a fine  powder  it  will  burn  when  ignited  in  the  air  to  OsO^.  Nitric 
acid  and  aqua  regia  convert  it  into  the  same  oxide.  The  compounds,  OsClj  and  OsO, 
OsClg  and  OsgO,,  OsOg  and  OsCl^,  are  very  similar  to  the  corresponding  ruthenium  deriva- 
tives. By  fusion  with  potassium  hydroxide  and  niter  we  get  potassium  osmate,  K.^OsO^, 
which  crystallizes  from  aqueous  solution  with  two  molecules  of  water  in  dark-violet  octa- 
hedra. The  most  stable  and  a very  characteristic  derivative  of  osmium  is  the  tetroxidcy 
OSO4,  which  is  produced  by  igniting  the  metal  in  the  air,  or  by  the  action  of  chlorine  on 
osmium  in  the  presence  of  water.  It  crystallizes  in  large  colorless  prisms,  which  fuse 
below  100°  and  distil  at  a somewhat  higher  temperature.  It  has  a very  sharp,  piercing 
odor  {bcfjL'^y  odor),  similar  to  that  of  sulphur  chloride.  It  dissolves  slowly  but  copiously 
in  water  ; the  solution  is  not  acid.  Reducing  and  organic  substances  precipitate  pulveru- 
lent osmium  from  it.  This  is  the  basis  of  its  application  in  microscopy. 

OsO^  and  RUO4  do  not  afford  corresponding  salts. 


RHODIUM  AND  IRIDIUM. 


Rh  = 103.0. 


Ir  193.0. 


These  metals  are  lighter  in  color  and  are  more  easily  fusible  than  ruthenium  and 
osmium.  (Iridium  fuses  at  1950°.)  When  pure  they  are  not  attacked  by  acids  or  aqua 
regia ; but  dissolve  in  the  latter  when  alloyed  with  platinum. 

Rhodiufn,  discovered  by  Wollaston  in  1803,  forms  three  oxides  : 


RhO,  RhgOg  and  RhOj, 


of  which  the  second  forms  salts  with  acids.  RhOj  results  when  rhodium  is  heated  with 
niter. 

Of  the  chlorides  only  RhCl,  is  known.  It  results  when  the  metal  is  heated  in  chlorine 
gas.  It  is  a brownish-red  mass.  It  forms  readily  crystallizing,  red-colored  (hence  its 
name,  from  podd£/f,  rose-red)  double  salts  with  alkaline  chlorides. 

Iridium y discovered  in  1804  by  Tennant,  has  perfectly  analogous  derivatives  : 

IrO,  IrjOg,  Ir02,  IrClg,  IrCl^. 

The  sesquichloride,  IrjClg,  formed  by  heating  iridium  in  chlorine,  is  an  olive-green, 
crystalline  mass,  insoluble  in  water  and  acids.  It  affbr&s  double  salts  with  the  alkaline 
chlorides,  e.g.  : 

KsIrClg  - 3H2O, 

which  crystallizes  from  water  in  green  or  brown  crystals.  They  are  also  produced  by 
the  action  of  sulphur  dioxide  upon  the  double  salts  of  hydrochloridic  acid. 

Iridium  Tetrachloride,  IrCl^,  is  produced  in  the  solution  of  iridium  or  its  oxide  in 
aqua  regia,  and  remains,  on  evaporation,  as  a black  mass,  readily  soluble  in  water  (with 
red  color)  to  hydrochloriridic  acid  : 

Il^IrClg. 

When  alkaline  chlorides  are  added  to  the  .solution  double  chlorides  are  precipitated,  e.g, : 

(NH,)2lrClg,  K^IrClg, 


412 


INORGANIC  CHEMISTRY. 


isomorphous  with  the  corresponding  double  chlorides  of  platinum.  When  a solution  of 
IrCl^  is  boiled  with  caustic  potash,  Ir(OII)^  will  be  precipitated,  d'he  name  iridium 
alludes  to  the  differently  colored  compounds  yielded  by  the  metal. 


PALLADIUM. 

I’d  = io6. 

Palladium,  in  addition  to  occurring  in  platinum  ores,  is  found  alloyed 
with  gold  (Brazil),  and  in  some  selenium  ores  (Hartz) ; it  has  a silver- 
white  color,  and  is  somewhat  more  fusible  (at  about  1400°)  than  platinum. 
When  finely  divided  it  dissolves  in  boiling  concentrated  hydrochloric, 
sulphuric,  and  nitric  acids.  When  ignited  in  the  air  it  at  first  becomes 
dull  by  oxidation,  but  at  a higher  temperature  the  surface  again  assumes 
a metallic  appearance.  It  was  discovered  in  1803  by  Wollaston  and 
named  after  the  planet  Pallas,  which  had  been  found  shortly  before. 

Palladium  absorbs  hydrogen  gas  (occlusion)  to  a much  greater  extent 
than  platinum  or  silver.  Freshly  ignited  palladium  leaf  absorbs  upwards 
of  370  volumes  of  hydrogen  at  ordinary  temperatures,  and  about  650 
volumes  at  90-100°.  A greater  absorption  may  be  effected  at  ordinary 
temperatures  in  the  following  manner: 

Water  is  decomposed  by  the  electric  current,  palladium  foil  being  used 
as  a negative  electrode.  The  liberated  hydrogen  is  then  taken  up  by  the 
palladium  (to  960  volumes) ; the  metal  expands  ( jlg-  its  volume),  becomes 
specifically  lighter,  but  retains  its  metallic  appearance  entire.  Accord- 
ing to  the  investigations  of  Debray,  the  compound  Pd^H2  is  produced, 
which  contains  dissolved  hydrogen,  and  deports  itself  similarly  to  an 
alloy.  Recent  investigations  show  that  palladium  hydride  is  a so-called 
“solid  solution”  of  hydrogen  in  palladium  (pp.  46,  256).  Palladium 
charged  with  hydrogen  usually  remains  unaltered  in  the  air,  and  in  a 
vacuum  ; it,  however,  sometimes  becomes  heated  in  the  air,  as  the  hydro- 
gen is  oxidized  to  water.  The  same  occurs  when  palladium  hydride  is 
heated  to  100°  ; in  vacuo,  all  the  hydrogen  escapes  as  gas.  Palladium 
hydride  is  a strong  reducing  agent,  like  nascent  hydrogen.  Ferric  salts 
are  reduced  to  the  ferrous  state;  chlorine  and  iodine  in  aqueous  solution 
are  converted  into  hydrochloric  and  hydriodic  acids. 

Palladium  black  absorbs  hydrogen  more  energetically  than  the  compact 
variety  (at  100°  upwards  of  980  volumes).  This  substance  is  obtained  by 
the  reduction  or  electrolysis  of  palladic  chloride.  If  palladium  sponge 
be  heated  in  the  air  until  the  white  metallic  color  becomes  black,  in  con- 
sequence of  the  superficial  oxidation,  it  will  absorb  hydrogen  very  ener- 
getically at  ordinary  temperatures,  and  partly  oxidize  it  to  water. 

When  ])alladium-sheet  or  sponge  is  introduced  into  the  flame  of  a 
spirit-lamp,  it  is  covered  with  soot;  this  is  due  to  the  fact  that  the 
metal  witlidraws  the  hydrogen  of  the  hydrocarbons,  and  carbon  is  set 
free. 

'I'here  are  two  series  of  ])alladium  compounds:  the  palladious,  PdX2, 
and  palladic,  PdX.^.  The  first  are  well  characterized  and  are  distinguished 
by  their  stability. 


PLATINUM. 


413 


Palladious  Chloride,  PdClj,  remains  as  a brown,  deliquescent  mass, 
on  evaporating  the  solution  of  j)alladium  in  aqua  regia.  It  yields  easily 
soluble  crystalline  double  salts,  with  alkaline  chlorides,  e.g.,  PdCl2.  2KCI 
= K^PdCl,. 

Palladious  Iodide,  Pdl2,  is  precipitated  from  palladious  solutions  by 
potassium  iodide  as  a black  mass,  insoluble  in  water.  As  palladium  chlo- 
ride and  bromide  are  very  soluble  in  water,  palladium  salts  can  be  used  to 
detect  iodine  in  the  presence  of  hydrobromic  or  hydrochloric  acid. 

Palladious  Oxide,  PdO,  is  a black  residue  left  upon  careful  ignition 
of  the  nitrate.  It  is  difficultly  soluble  in  acids.  When  heated,  it  loses 
oxygen,  and  forms  palladium  suboxide,  Pd20. 

When  palladium  is  dissolved  in  sulphuric  or  nitric  acids,  the  corre- 
sponding salts  are  produced. 

The  sulphate,  PdSO^  -f  2H2O,  is  composed  of  brown  crystals,  readily 
soluble  in  water.  Much  of  the  latter  decomposes  it. 

Palladio  Chloride,  PdCl^,  Hydrochlorpalladic  Acid,  H2PdClg,  is 
formed  when  the  metal  is  dissolved  in  aqua  regia.  They  decompose,  on 
evaporation,  into  PdCl2  and  CI2.  When  potassium  or  ammonium  chlo- 
ride is  added  to  their  solutions,  red-colored  difficultly  soluble  double 
chlorides  crystallize  out ; they  are  analogous  to  the  corresponding  salts 
of  platinum. 


PLATINUM. 

pt  = 194.8. 

The  separation  of  platinum  from  the  ore  was  described  on  p.  409. 
(See  Myliusand  Foerster,  Ber.  25  (1893),  665.)  The  metal  has  a grayish- 
white  color,  and  a specific  gravity  of  21.4.  It  is  very  tough  and  malleable, 
and  may  be  drawn  out  into  very  fine  wire  and  rolled  into  foil.  At  a high 
heat  it  softens  without  melting,  and  may  be  easily  welded.  It  fuses  in 
the  oxy hydrogen  flame  (at  about  1770° — Violle),  and  is  somewhat  volatile. 
On  fusion,  it  absorbs  oxygen,  which  it  gives  up  again  on  cooling  (like 
silver).  At  ordinary  temperatures,  it  also  condenses  hydrogen  and 
oxygen  upon  its  surface;  as  foil  and  sponge,  but  very  few  volumes ; as 
platinum  black,  about  100  times  its  volume  of  oxygen  and  310  times  its 
volume  of  hydrogen.  Two  hundred  volumes  of  the  latter  form  water 
with  the  oxygen  always  present  in  the  sponge  (Z.  f.  anorg.  Ch.  10  (1895), 
178).  These  gases  are  fully  expelled  at  a red  heat.  Platinum  sponge  is 
obtained  as  a fine  black  powder,  if  reducing  substances,  like  zinc,  be  added 
to  solutions  of  platinic  chloride  or  upon  boiling  the  solution  with  sugar 
and  sodium  carbonate;  it  remains  on  the  ignition  of  PtCl^.  2NH^C1.  The 
production  of  various  reactions  is  due  to  the  power  of  platinum  to  con- 
dense oxygen  ; thus  hydrogen  will  inflame  in  the  air,  if  it  be  conducted 
upon  platinum  sponge ; sulphur  dioxide  combines  with  oxygen  at  100° 
to  form  the  trioxide.  At  a red  heat  platinum  permits  free  passage  to 
hydrogen,  while  it  is  not  permeable  by  oxygen  and  other  gases  (j).  93). 

Platinum  is  not  attacked  by  acids;  it  is  only  soluble  in  liquids  gener- 
ating free  chlorine,  e.  g.,  aqua  regia.  In  consequence  of  this  resistance 


414 


INOROANIC  CHEMISTRY. 


to  acids,  and  its  nnalterability  upon  ignition,  tliis  metal  answers  as  an 
imdecomposable  material  for  the  production  of  chemical  crucibles,  dishes, 
wire,  etc.  The  usual  presence  of  iridium  in  ordinary  platinum  increases 
its  durability. 

The  alkaline  hydroxides,  sulphides,  and  cyanides  attack  it  strongly  at 
a red  heat.  It  forms  readily  fusible  alloys  with  ])hosphorus,  arsenic,  and 
many  heavy  metals,  especially  lead,  and  many  heavy  metals  are  reduced 
from  their  salts  by  platinum.  Therefore  such  substances  must  not  be 
ignited  in  platinum  crucibles,  etc. 

Platinum,  like  palladium,  forms  platifious,  PtX.^,  and  platinic,  PtX^, 
derivatives  ; in  the  first  it  is  more  basic,  in  the  latter  more  acidic. 

Platinic  Chloride,  PtCl^,  is  obtained  by  the  solution  of  platinum  in 
aqua  regia,  and  when  the  liquid  is  evaporated  with  an  excess  of  hydro- 
chloric acid,  hydrochlorplati ri ate , 

PtCh.2HCl  ^ HjPtCle, 

crystallizes  with  six  molecules  of  water  in  brownish-red,  deliquescent 
prisms.  It  forms  characteristic  double  chlorides, 

PtCh.2KCl, 

with  ammonium  and  potassium  chlorides.  They  are  the  chlorplatinaies. 
Formerly  they  were  considered  double  salts.  They  are  difficultly  soluble 
in  water;  hence,  on  mixing  the  solutions,  they  immediately  separate  out 
as  a crystalline  yellow  powder.  Ignition  completely  decomposes  the 
ammonium  salt,  leaving  spongy  platinum. 

Platinum  chloride  yields  similar  insoluble  salts  with  the  chlorides  of 
rubidium,  caesium,  and  thallium,  while  that  with  sodium, 

NajPtCle  + 6H.,0, 

is  very  soluble  in  water.  Platinum  may  be  very  readily  separated  from 
other  metals  by  recrystallizing  the  sodium  salt  from  hot  water  rendered 
alkaline  by  soda  (Finkener). 

On  adding  sodium  hydroxide  to  platinic  chloride  and  then  supersatu- 
rating with  acetic  acid,  there  separates  a reddish-brown  precipitate  of 
platinic  hydroxide,  Pt(OH)^.  This  dissolves  readily  in  acids  (excepting 
acetic),  with  formation  of  salts.  The  oxygen  salts,  as  Pt(S04")2,  are  very 
unstable.  The  hydroxide  has  also  an  acidic  character  (^platinic  acid), 
and  dissolves  in  alkalies,  yielding  salts  with  them.  These,  also,  result 
on  fusing  platinum  with  potassium  and  sodium  hydroxide.  The  barium 
salt, 

is  ])recipitated  from  platinic  chloride,  by  barium  hydroxide,  as  a yellow, 
crystalline  compound.  The  acidic  nature  of  its  hydroxide  allies  platinum 
to  gold.  If  hydrogen  sulphide  be  conducted  through  platinic  solutions, 
black  platinum  disulphide,  PtSj,  is  precipitated;  it  is  soluble  in  alkaline 
siil|)hides,  with  formation  of  sulpho-salts. 

Platinous  Chloride,  PtCl.^,  is  a green  powder,  insoluble  in  water. 


SPECTRUM  ANALYSIS.  415 

remaining  after  heating  platinic  chloride  to  200°.  It  forms  double  salts 
with  alkaline  chlorides,  e.  g.  : 

PtCl, . 2NaCl. 

When  digested  with  potassium  hydroxide  it  yields  the  hydroxide, 
Pt(OH),. 

Cyanogen  Compounds. — Like  cobalt,  platinum  forms  double  cy- 
anides corresponding  to  the  ferrocyanides.  When  platinous  chloride  is 
dissolved  in  potassium  cyanide  platinum  potassium  cyanide,  K.2Pt(CN)^  -|- 
4H2O,  crystallizes  on  evaporation  in  large  prisms  exhibiting  magnificent 
dichroism ; in  transmitted  light  they  are  yellow  and  in  reflected  light 
blue.  This  salt  must  be  viewed  as  the  potassium  compound  of  hydro- 
pi  atino- cyanic  acid,  H2Pt(CN)4.  When  separated  from  its  salts  it  crys- 
tallizes in  gold-yellow  needles.  Its  salts  with  heavy  metals  are  obtained 
by  double  decomposition,  and  all  show  a beautiful  play  of  colors.  The 
barium  salt,  barium  platino-cyanide,  BaPt(CN)4  4H2O,  is  used  to 
render  the  Rontgen  rays  visible. 


Platinum-ammonium  Compounds. — There  is  a whole  series  of 
these,  which  must  be  viewed  as  platinum  bases  and  their  salts.  They  are 
constituted  according  to  the  following  empirical  formulas  : 

{a)  Platosamines  (platinous  series)  : 

PtR2(NH3)2;  PtR2(NH3)3 ; PtR2(NH3b. 

[{h)  Platinamines  (platinic  series)  : 

PtR,(NH3);  PtR,(NH3)2;  PtR,(NH3)3 ; PtR,(NH3),. 

(R  = OH,  Cl,  Br,  I,  NO3,  etc.). 

Diplatosamine  derivatives,  <?.  g.,  (NH3)^Pt2R2,  are  also  known.  They 
evidently  contain  radicals  or  ions,  consisting  of  platinum,  the  group  NH3, 
and  halogens  (also  NO3,  etc.).  They  result  from  the  action  of  ammonia 
upon  platinous  chloride.  When  the  acid  residues  are  replaced  by 
hydroxyl  groups  platinum  bases  are  formed,  e.  g.,  Pt(NH3)^(OH)2,  which 
resemble  alkaline  hydroxides  in  their  chemical  properties.  The  other 
platinum  metals  form  similar  amine  derivatives.  The  nature  and  chemical 
constitution  of  these  interesting  compounds  is,  however,  not  fully  ex- 
plained. 


SPECTRUM  ANALYSIS.* 


We  observed  that  many  substances,  when  introduced  into  a non-lumi- 
nous  flame,  imparted  to  it  a characteristic  coloration.  Thus,  the  sodium 


* A more  exhaustive,  concise,  and  distinct  presentation  of  the  s])ectrum  plienomena  may 
be  found  in  Herman  VV.  Vogel’s  “ Prakti.sche  .Spektralanalyse  irdi.scher  Stoffe,”  1889,  and 
in  “ Lehrbuch  der  .Spektralanalyse,”  by  H.  Kayser,  1883. 


4i6 


INORGANIC  ClIKMIS'I'kV. 


compounds  color  it  yellow;  the  iiotassiiim,  violet;  thallium,  green,  etc. 
The  decomj)osition  of  the  light  thus  obtained,  and,  indeed,  of  every  liglu, 
by  means  of  a prism,  and  the  study  of  the  resulting  spectrum,  form  the 
basis  of  spectrum  analysis,  established  in  1859  by  Kirchhoff  and  bunsen. 

Every  substance,  solid  or  liquid,  heated  to  white  heat  {e.  g.,  molten 
platinum;  lime  heated  in  the  oxyhydrogen  flame;  the  ordinary  flame 
containing  glowing  particles  of  carbon)  emits  rays  of  every  refrangi- 
bility ; and  hence  furnishes  a cotitinuous  speci7'uin,  which  firings  to  view 
all  the  colors  of  the  rainbow,  from  red  to  violet,  if  the  light  be  conducted 
through  a prism.  Glowing  gases  and  vajiors,  on  the  contrary,  whose 
molecules  can  execute  unobstructed  oscillations,  emit  light  of  definite 
refrangibility,  and,  therefore,  give  spectra,  consisting  of  single,  bright 
lines.  Thus,  the  spectrum  of  the  yellow  sodium  flame  is  recognized  as 


composed  of  one  very  bright  yellow  line,  which  by  increased  magnifying 
power  is  shown  to  consist  of  two  lines  lying  very  near  each  other.  The 
violet  potassium  light  gives  a spectrum,  consisting  of  a red  and  a blue 
line.  The  crimson  strontium  light  shows  in  the  spectrum  several  dis- 
tinct red  lines  and  a blue  line.  (See  the  Spectrum  Plate.)  Each  of  these 
lines  corresjionds  to  rays  of  definite  wave-length,  which  are  differently 
refracted  from  all  the  others;  hence  the  lines  appear  in  the  same  spec- 
trum at  the  same  point,  when  rays  of  their  wave-length  fall  on  the  slit. 

If  substances  affording  different  colors  be  introduced  into  a flame,  the 
most  intense  color  generally  obscures  the  others  ; in  the  spectrum,  how- 
ever, each  individual  substance  shows  its  peculiar  bright  lines,  which 
ai)pcar  simultaneously  or  succeed  each  other,  according  to  the  volatility 
of  the  various  substances. 


SPECTRUM  ANALYSIS. 


417 


The  spectrum  apparatus  or  spectroscope,  shown  in  Fig.  69,  serves  for 
the  observation  of  the  spectra. 

In  the  middle  of  the  apparatus  is  a flint-glass  prism  P.  At  the  further 
end  of  the  tube  A is  a movable  vertical  slit,  in  front  of  which  is  placed 
the  light  to  be  investigated.  The  entering  light  rays  are  directed  by  a 
collecting  lens  into  the  tube  A,  upon  the  prism,  and  the  refracted  rays 
(the  spectrum)  are  observed  by  the  telescope  B.  The  tube  C is  employed 
to  ascertain  the  relative  position  of  the  spectrum  lines.  This  is  provided 
at  the  outer  end  of  the  tube  S with  a transparent  horizontal  scale.  When 
a luminous  flame  is  placed  before  the  scale  its  divisions  are  reflected  from 
the  prism  surface  and  thrown  into  the  telescope  B.  We  then  see  the 
spectrum  to  be  studied  and  the  scale  divisions  in  B at  the  same  time,  and 
can  readily  determine  the  relative  position  of  the  lines  of  the  spectrum. 
To  study  two  spectra  at  the  same  time,  and  compare  them,  a three-sided, 
right-angled  glass  prism  is  attached  in  front  of  one-half  (the  lower  or 
upper)  of  the  slit  of  the  tube  A ; this  directs  the  rays  of  a light  placed 
at  the  side  (/,  Fig.  69)  through  A upon  the  prism  P.  By  means  of  B, 
two  horizontal  spectra  will  be  observed,  one  above  the  other,  and  be- 
tween are  the  bright  divisions  of  the  scale. 

Adjustment  of  the  Spectroscope. — To  observe  the  spectra  in  the  apparatus  described,  it 
is  necessary  first  to  adjust  the  same  correctly.  The  tube  A (objective  or  collimator  tube) 
contains,  besides  the  slit,  also  a lens  (collimator  lens),  which  serves  to  render  parallel  the 
bunch  of  rays  proceeding  from  the  slit ; hence,  the  latter  must  be  accurately  placed  in 
the  focus  of  the  lens.  This  is  best  accomplished  as  follows  : The  telescope  {B)  is  drawn 
out  and  adjusted  for  some  distant  object,  that  it  may  be  adapted  for  the  reception  of 
parallel  rays  ; observe  the  slit  illuminated  by  the  sodium  chloride  flame,  and  change  its 
position  with  reference  to  the  collimator  lens  until  its  image  appears  sharply  defined  in 
the  telescope  B.  The  scale  in  tube  .S'  is  similarly  adjusted.  To  have  the  spectrum  lines 
sharply  defined,  the  slit  must  be  made  quite  narrow  ; for  feebly  luminous  lines  it  must, 
however,  be  widened. 

The  determination  of  the  position  of  the  lines  of  the  spectra  is  usually  effected  by 
means  of  a scale.  Since  the  refraction  and  dispersion  of  the  rays  and  therefore  the  ex- 
pansion of  the  spectrum  and  the  length-relations  of  its  variously  colored  parts  are  influ- 
enced by  the  angle  and  the  quality  of  the  glass  of  the  prism,  the  scale  indications  of 
different  forms  of  apparatus  are  not  directly  comparable.  To  attain  this  it  is  ordinarily 
sufficient  to  indicate  how  the  most  important  Fraunhofer  lines  (see  below)  are  divided 
upon  the  scale. 

The  spectra  represented  on  the  plate  are  then  produced  by  the  fre- 
quently mentioned  colors  shown  by  compounds  of  the  alkali  and  alkaline 
earth  metals  when  heated  in  the  non-luminous  flame  of  the  Bunsen  burner. 

The  other  metals,  however,  require  a much  higher  temperature  than 
that  of  the  gas  flame  for  their  conversion  into  gases.  To  vaporize  them 
and  observe  their  spectra,  the  electric  spark  is  made  to  pass  from  elec- 
trodes constructed  of  them.  A very  convenient  and  universally  appli- 
cable procedure  consists  in  passing  the  sparks  between  the  surface  of  the 
solution  of  the  salt,  into  which  the  negative  pole  dips,  and  a platinum 
wure  (positive  pole)  placed  before  the  slit  of  the  spectroscope  [R.  Bun- 
sen, Pogg.  Ann.  (1873)  I55»  230,  336;  Z.  f.  anal.  Ch.  (1876)  15,  68, 
and  M.  Lecoq  de  Boisbaudran,  S])ectres  lumineux,  Paris,  1874].  Many 
of  the  metal  spectra,  obtained  in  this  way,  are  remarkably  rich  in  lines. 
Thus,  over  450  lines  have  been  established  for  iron. 


4i8 


ORGANIC  CHRMISIRY. 


Tlie  speclra  of  the  gases  may  l)e  obtained  l)y  means  of  llie  ('.cissler 
tubes.  Tlie  latter  contain  tlie  gas  in  a rarefied  condition  and  they  give 
their  own  i)eculiar  colors  on  jiassing  induction  sparks  through  tliem. 
Hydrogen  gives  a red  light ; its  spectrum  consists  of  a bright  red,  a bine, 
and  a green  line.  Nitrogen  gives  a violet  light  and  affords  a s[)ectrum 
of  many  lines  and  bands. 

These  methods  give  ns  a means  for  readily  distinguishing  the  individual 
chemical  elements,  and  e\‘en  detecting  them  in  traces.  Since  the  year 
i860  various  new  elements,  <f.  , cmsinm,  rubidium,  thallium,  indium, 

scandium,  gallium,  germanium,  and  several  others,  not  accurately  studied, 
have  been  discovered  by  their  aid.  Argon,  helium  and  the  other  sub- 
stances discovered  recently  in  the  air  are  esi)ecially  distinguished  by  their 
spectra. 

In  addition  to  the  bright  emissioti  spectra  just  described  there  are  yet 
dark  absorption  spectra.  If  a white  light  giving  an  uninterru|)ted  spec- 
trum be  allowed  to  pass  through  different  transparent  bodies,  the  latter 
will  absorb  rays  of  definite  refrangibility,  allowing  all  others  to  ])ass. 
Therefore  we  observe  the  sun  spectrum  interrupted  by  dark  lines  or  bands 
in  the  spectroscope.  The  solutions  of  didymium  and  erbium  absorb  cer- 
tain rays,  and  show  corresponding  dark  lines  in  the  spectrum.  The 
gases  deport  themselves  similarly.  White  light  that  has  traversed  a broad 
layer  of  air  shows  several  dark  lines  in  the  spectrum;  the  air  absorbed 
the  corresponding  rays.  This  power  of  absorption  is  possessed  to  a much 
higher  degree  by  all  incandescent  gases  or  vapors.  If  a white  light,  like 
the  Drummond  calcium  light,  be  conducted  through  the  yellow  sodium 
flame  (through  glowing  sodium  vapors),  a dark  line  will  appear  in  the 
spectrum  of  the  white  light,  and  its  position  will  correspond  exactly  to 
that  of  the  yellow  sodium  line;  the  latter  thus  appears  converted  into  a 
dark  line.  Such  spectra  are  designated  the  inverted  spectra  of  the  corre- 
sponding metals.  The  inverted  spectra  of  all  elements  may  be  obtained 
in  this  way,  and  they  correspond  accurately  to  the  direct  bright  spectra. 
The  cause  of  these  phenomena  lies  in  the  proposition  deduced  by  Kirch- 
hoff  from  the  undulatory  theory  of  light,  that  the  ratio  betiveen  the  emis- 
sive a7id  absorptive  potver  is  the  same  for  almost  all  bodies  at  like  tempera- 
tures. According  to  this  incandescent  gases  only  absorb  rays  of  just  the 
same  refrangibility  as  those  which  they  emit.  For  example,  when  bright 
white  light  is  passed  through  the  yellow  sodium  flame  the  yellow  rays  of 
the  former  are  absorbed  and  retained,  while  all  others  pass  on  almost 
entirely  unaltered.  Therefore,  in  the  rainbow  spectrum  of  white  light 
the  yellow  rays  of  definite  refrangibility  will  be  absent ; and  if  the  other 
refracted  rays  of  the  white  light  are  brighter  than  the  yellow  rays  emitted 
from  the  sodium  flame,  the  latter  will  be  relatively  darker ; a dark  line 
will  therefore  make  its  appearance. 

These  jihenomena  have  presented  a new  and  wide  province  to  spectral 
analysis,  inasmuch  as  they  open  up  avenues  for  the  investigation  of  the 
chemical  nature  of  the  sun  and  other  bodies. 

It  is  knf)wn  that  the  bright  rainbow  sun  spectrum  is  intersected  by  a 
number  of  dark  lines  which  have  been  called  the  Fraunhofer  lines  from 
their  discoverer.  Kirchhoff  has  shown  that  these  lines  can  be  easily  ac- 


SPECTRUM  ANALYSIS. 


419 


counted  for,  after  what  has  already  been  said,  by  the  following  hypothesis 
upon  the  nature  of  the  sun  ; The  latter  consists  of  a solid  or  liquid  lumi- 
nous nucleus,  surrounded  by  an  atmosphere  of  incandescent  gases  and 
vapors.  Then  the  uninterrupted  spectrum  of  the  glowing  nucleus  must 
be  intersected  by  the  dark  lines  of  the  inverted  spectra  of  those  gases  and 
vapors  which  occur  in  the  sun’s  atmosphere.  An  accurate  comparison  of 
the  Fraunhofer  lines  with  the  spectrum  lines  of  the  various  elements  has 
revealed  the  fact  that  iron,  sodium,  magnesium,  calcium,  chromium, 
nickel,  barium,  copper,  zinc,  and  hydrogen  are  present  in  the  sun’s 
atmosphere.  Tims  dark  lines  have  been  found  in  the  sun’s  spectrum 
corresponding  to  all  of  the  450  lines  of  the  iron  spectrum.  According 
to  Rowland  34  of  the  elements  known  to  us  are  present  in  the  sun ; 15 
are  absent,  and  much  doubt  prevails  as  to  the  presence  of  the  remainder. 
Helium,  first  discovered  in  the  earth  in  1895,  was  known  to  exist  in  the 
sun’s  photosphere  since  1868.  The  inferences  upon  the  chemical  consti- 
tution of  the  sun  possess  as  much  and,  indeed,  a higher  degree  of  proba- 
bility than  falls  to  many  other  deductions. 

All  the  fixed  stars  thus  far  investigated  possess  a constitution  like  that 
of  the  sun.  They  give  spectra  intersected  by  dark  lines,  and  therefore 
consist  of  incandescent  nuclei  surrounded  by  gaseous  atmospheres.  The 
spectra  of  nebulae,  however,  only  show  bright  lines ; hence  these  consist 
of  uncondensed,  incandescent  masses  of  vapor. 


INDEX 


A. 

Absorptiometer,  120 
Acetylene,  15 1,  154 
Acids,  39,  60,  84 
haloid,  60 
meta-,  201 
normal,  201 
ortho-,  201 
oxygen,  6 1 
radicals,  1 74 
Active  oxygen,  84 
chlorine,  304 
Affinity,  chemical,  23 
Air,  composition  of,  118,  119,  1 20,  I21, 
122,  123 
destroyed,  118 
fire,  1 18 

liquefaction  of,  1 17 
liquid,  1 17 
Algaroth,  1 47 
Alkali  metals,  272 
Alkalies,  272 
Alkaline  earths,  300 
recognition  of,  310 
Allotropy,  87 
Alloys,  255 

Alum,  ammonium,  353 
burnt,  353 
cubical,  353 
ordinary,  352 
Alums,  352 
Aluminates,  350 
Aluminium,  346 

bronze,  336,  347 
chloride,  348 
hydroxides,  350 
oxide,  349 
phosphate,  353 
silicates,  353 
sulphate,  352 
Amalgams,  257 
Ammonia,  126 

chemical  properties  of,  128 
physical  properties  of,  127 
quantitative  composition  of,  129 
thermo-chemical  deportment  of,  1 29 


Ammonia-cobalt  compounds,  407 
Ammonium,  129 
amalgam,  296 
carbonate,  298 
chloride,  297 
compounds,  129,  296 
hyponitrite,  297 
nitrate,  297 
nitride,  298 
nitrite,  297 
persulphate,  297 
phosphates,  298 
sulphate,  297 
sulphides,  299 
Ampere,  law  of,  75 
Anglesite,  371 
Anhydrite,  304 
Anhydrides  of  acids,  172 
Animal  charcoal,  15 1 
Anion,  92 
Anode,  92 
Anthracite,  15 1 
Antichlor,  289 
Antimony,  146 
acids,  224 
bromide,  148 
butter  of,  147 
chlorides,  147,  148 
cinnabar,  225 
fluorides,  148 
iodides,  148 
mirror,  147 
oxides,  223,  224 
oxychloride,  147 
spots,  147 
sulphides,  225 
Antimonyl,  224 
Antozone,  88 
Apatite,  304 
Apparatus  of  Carr^,  127 
Marsh,  144 
Aqua  regia,  203 
Aragonite,  305 
Argentan,  336,  405 
I Argon,  1 16,  124 
I Argyroditc,  362 

421 


422 


INDEX. 


Arsenic,  143 

acids,  221,  222 
bloom,  220 
bromide,  146 
chloride,  146 
fluoride,  146 
iodides,  146 
Marsh’s  test  for,  145 
mirror,  144 
pentoxide,  222 
spots,  144 
sulphides,  222,  223 
sulpho-salts,  223 
trioxide,  220 
Arsenious  acid,  221 
Arsenites,  221 
Arsine,  144,  I45 
Artificial  cinnabar,  327 
Atmosphere,  1 16 

gases  discovered  in,  123,  124 
Atomic  compounds,  170 
heat,  253 
hypothesis,  69,  72 
-molecular  theory,  73 
volume,  252 
weight,  167 

choice  of,  69 
relative,  25,  72 
table  of,  26 
thermal,  254 
Atomicity,  165 
Atoms,  24,  75,  167 
Auric  acid,  344 

compounds,  344,  345 
Auripigment,  223 
Aurous  compounds,  343 
Avogadro,  law  of,  75,  98 
Azoimide,  132 
Azote,  118 
Azurite,  335 

B. 

Barium,  308 

carbonate,  310 
nitrate,  309 
oxide,  308 
peroxide,  309 
persulphate,  309  • 
sulphate,  309 
sulphide,  310 
T^aryta  water,  309 
Bases,  39,  60,  84 
basicity  of  acids,  1 72,  214 
Bauxite,  353 
Beryllium,  317,  318 
l>csscmer  steel,  396 
Bismuth,  148,  372,  373 
alloys,  373 
miiuic,  373 


Bismuth,  oxide,  373 
oxychloride,  372 
subnitrate,  373 
Bismuthic  acid,  373 
Bleaching,  53 
lime,  303 
Borax,  243,  294 
Boric  acid,  242 

meta,  242 

Boro- fluoride,  hydrogen,  242 
l)oron,  241,  294 
carbide,  243 
chloride,  241 
fluoride,  242 
hydride,  241 
nitride,  243 
trioxide,  242 
Brass,  335 
Britannia  metal,  364 
Bromic  acid,  179 
Bromine,  53 
solid,  54 
Bronzes,  336 

C. 

Cadmium,  320 
oxide,  320 
sulphate,  320 
sulphide,  321 
Cadmous  oxide,  320 
Caesium,  284 
Calamine,  319 
Calcite,  305 
Calcium,  300 
carbide,  307 
carbonate,  305 
chloride,  302 
fluoride,  302 
hydride,  301 
hydroxide,  301 
hypochlorite,  303 
nitrate,  304 
nitride,  301 
oxide,  301 
peroxide,  302 
phosphates,  304,  305 
silicate,  306 
sulphate,  304 
sulphides,  307 
Calomel,  323 
Calorie,  22,  66,  253 
Caput  mortuum,  194,  399 
Carbides,  metallic,  257 
Carbon,  149 

amorjflious,  151 
chlorides,  159 

compounds  with  hydnigen,  151 
dioxide,  227,  228,  229,  230 
liquid,  228 


INDEX. 


423 


Carbon  bisulphide,  234 
gas,  15 1 
group,  149 

monoxide,  231,  232,  233 
oxysulphide,  235 
Carbonates,  23 1 
Carbonic  acid,  230,  231 

amido-derivatives  of,  234 
Carbonyl  chloride,  234 
Carborundum,  163 
Carry’s  ice  machine,  127 
Carnallite,  276,  315 
Cassiterite,  363 
Catalytic  reactions,  103 
Caustic  potash,  275 
soda,  286 
Cement,  302 
Cementation  steel,  396 
Cerium,  355 
Cerrusite,  367,  371 
Chalk,  305 
Chamber  acid,  19 1 
Chameleon  minerale,  39 1 
Charcoal,  15 1 
animal,  151 
Chemical  affinity,  23 
combination,  19 
constitution,  168 
elements,  19 
energy,  23 
equations,  27 
formulas,  25,  27 
industries,  importance  of,  1 8 
phenomena,  19 
structure,  168 
symbols,  25 
tension,  23 

Chlor-detonating  gas,  57 
Chloric  acid,  177 
Chloride  of  lime,  303 
Chlorine,  49 

dioxide,  176 
hydrate,  52 
preparation,  49 
properties,  51,  52 
tetroxide,  176 
thionyl,  185 
trioxide,  176 
Chlorites,  176 
Chlorous  acid,  176 
Chlorsulphonic  acid,  195 
Chromates  of  potassium,  379 
Chrome  yellow,  380 
Chromic  acid,  377 

anhydride,  378 
chloranhydrides,  380 
compounds,  375 
oxide,  376 
Chromium,  374 


Chromium  group,  373,  374 
Chromous  compounds,  375 
Chromyl  chloride,  380 
Cinnabar,  322,  327 
9^y,  353 

Coal,  anthracite,  151 
bituminous,  15 1 
Cobalt,  406 

ammonia-compounds,  407,  408 
cyanides,  408 
Cobaltic  compounds,  407 
CobaltoLis  compounds,  406 
ultramarine,  407 
Coke,  15 1 
Colcothar,  399 
Colloids,  238 
Combining  weight,  70 
weights,  law  of,  72 
Condensation  of  gases,  47 
Conservation  of  energy,  2I 
Constant  proportions,  law  of,  69,  72 
Constantan,  405 
Copper,  330,  331,  332 
alloys,  335 
carbonates,  335 
hydride,  333 
sulphate,  334 
Corrosive  sublimate,  325 
Corundum,  346,  349 
Critical  condition,  47 
pressure,  47,  229 
temperature,  47 
volume,  47 
Crystallography,  31 
Crystalloids,  238 
Cupric  compounds,  334 
Cuprite,  332 

Cuprous  compounds,  332 
Cyanogen,  235 

compounds,  235 

derivatives  of  iron,  401,  402,  403 


D. 

Davy’s  lamp,  59 
Decipium,  27,  356 
Density,  gas,  74,  79,  80 
Determined  compounds,  c2 
Dialysis,  237 
Diamide,  131 

Diammonium  compounds,  310,  311,  312 
Diamond,  1 50 
Didymium,  355 
Dissociation,  94 
Dithionic  acid,  198 
Doctrine  of  linking  of  atoms,  172 
I )oberciner’s  lamp,  46 
Drummond  light,  83 
‘ Dysprosium,  27,  356 


424 


INDEX. 


E. 

Earth  metals,  345,  346,  354 
Eau  de  Javelle,  278 

Labarraque,  278 

Electro-chemical  equivalents,  265 
Electrolysis  of  salts,  262 
Electrolytic  dissociation  theory,  269 
I^lements,  chemical,  19 
Endothermic  compounds,  30 
reactions,  30,  67 
Energy,  22,  29 
chemical,  22 
degradation  of,  28 
electric,  22 
mechanical,  22 
radiant,  22- 
Entropy,  23 
Equivalence,  167 
Equivalent  weight,  167 
Erbium,  356 
Ethane,  153 
Ethylene,  153 

tetrachloride,  159 
Eudiometer,  120 
Eudiometric  analysis,  120 
Exothermic  compounds,  30 
reactions,  30,  67 


E. 

Faraday’s  law,  263 
Ferric  acid,  401 
chloride,  400 
compounds,  399,  400 
Ferrous  compounds,  397,  398 
sulphate,  398 
sulphide,  399 
Fire-damp,  152 
Fischer’s  salt,  407 

Flame,  nature  of,  155,  156,  157,  158 
oxidizing,  158 
reducing,  158 
Fluorine,  56 
Fluorite,  30 1 
Formulas,  structural,  168 
Fulminating  gold,  348 
silver,  338 

Furnace,  electric,  253 

G. 

Gadolinite,  354 
Gadolinium,  27,  356 
Gahnite,  351 
Galenite,  367 
Gallium,  358 


Gallium  compounds,  358 
group,  355 
(ias,  laughing,  21 1 

condensation  of,  47 
critical  condition  of,  47 
pressure  of,  47 
volume  of,  47 
density  of,  79 
Gases,  diffusion  of,  123 

drying  and  purifying,  42 
general  properties  of,  73 
measuring  of,  120 
General  sciences,  17 
Germanium,  164,  238,  362 
Germanic  compounds,  363 
Germanous  compounds,  362 
Glass,  306 
Glauber’s  salt,  287 
Glover  tower,  190 
Glucinum,  317 
Gold,  342,  343 

chlorides,  343,  344 
fulminating,  345 
oxides,  344 
Graphite,  150 
Graphitite,  150 
Green  cinnabar,  407 
Greenockite,  321 
Groups,  167 
Guignet’s  green,  376 
Gunpowder,  280 
Gypsum,  304 


H. 

Halogen  compounds  of  metals,  257 
hydrides,  57 
Halogenides,  92 
Halogens,  49 

compounds  with  one  another,  68 
. oxygen  compounds  of,  173 
thermo-chemical  deportment  of,  66 
Haloid  acids,  57,  61 
salts,  61 

Heat,  atomic,  253 
latent,  90 
modulus,  31,  66 
specific,  253 
unit  of,  66,  89 
Helium,  116,  125 
Hepar,  283 
Ilolmium,  27,  356 
Horn-silver,  339 
Hydrates,  258 
Hydraulic  cement,  302 
Hydrazine,  131 
hydrates,  31 1 
* Hydrazoic  acid,  132 


INDEX. 


425 


Hydriodic  acid,  63 
Hydrobroraic  acid,  62 
Hydrocarbons,  saturated,  153 
unsaturated,  153 
Hydrochloric  acid,  59 
Hydrochlorpalladic  aoid,  413 
Hydrocyanic  acid,  236 
Hydrofluoric  acid,  65 
Hydrosulphuric  acid,  109 
Hydrosulphurous  acid,  86 
Hydrogen,  40 

antimonide,  147 
bromide,  61 
chloride,  57 

chemical  properties  of,  60 
cyanide,  236 
fluoride,  64 
iodide,  63 
pentasulphide,  iio 
peroxide,  99 

detection  of,  103 
properties  of,  loi 
thermo-chemistry  of,  103 
persulphide,  iio 
preparation,  40 
properties,  43,  45 
selenide,  1 12 
silicofluoride,  163 
sulphide,  107,  108 

molecular  formula  of,  109 
thermo-chemical  deportment  of, 
109 

telluride,  113 
Hydroxides,  84 
Hydroxyl,  104 
Hydroxylamine,  130,  131 
Hypobromous  acid,  179 
Hypochlorous  acid,  1 74 
oxide,  174 
Hyponitric  acid,  206 
Hyponitrous  acid,  212 
Hypophosphoric  acid,  216 
Hypophosphorous  acid,  214 
Hypophosphites,  215 
Hyposulphites,  197 
Hyposulphurous  acid,  197 


I. 

Illuminating  gas,  155 

Indestructibility  of  matter,  principle  of,  20 
Indium,  358 

compounds,  359 
Iodic  acid,  180 

anhydride,  l8l 
Iodine,  54,  55 
bromide,  69 
chloride,  68 

36 


Iodine  fluoride,  69 
trichloride,  69 
lodonium  hydroxide,  1 80 
Iridium,  41 1 
Iron,  393,  394 
alums,  400 
carbide,  394,  402 
carbonyl,  404 
cast,  394 
disulphide,  400 
fluid,  394 
group,  392 
metallurgy,  395,  396 
pentacarbonyl,  404 
technical,  394 
tetracarbonyl,  404 
vitriol,  398 
weld,  394 
wrought,  394 
Isomerism,  87 
Isomorphism,  255 


— V K. 

Kaolin,  353 
Kathode,  92 
Kation,  92 
Kelp,  54 

Kermes  mineral,  225 
Kieserite,  287,  315 
Krypton,  1 16 


L. 


Lanthanum,  355 
Lapis  lazuli,  354 
Latent  heat,  90 
Laughing  gas,  21 1 
Law  of  Boyle,  74 

Dalton,  74 

definite  proportions,  69 
Faraday,  263 
Gay-Lussac,  74 
Mariotte,  74 
Lead,  164,  238,  364 
alloys,  368 
carbonate,  37 1 
chamber  crystals,  191 
chloride,  370 
chromate,  380 
iodide,  370 
nitrate,  37 1 
oxides,  369 
peroxide,  369 
red,  369 
sulphate,  371 
sulphide,  371 
tetrabromide,  371 
tetrachloride,  369,  370 


INDEX. 


426 


Lead  telra-iodide,  37 1 
tree,  368 
wliite,  371 
Lignite,  151 
Lime,  301 

cliloride  of,  303 
light,  83 
Litharge,  369 
Lithium,  295 

carbonate,  296 
chloride,  296 
hydride,  295 
phosphate,  296 
Lunar  caustic,  341 


M. 

Magnesia,  314 
usta,  314 
Magnesium,  313 

group,  metals  of,  312 
carbonate,  316 
chloride,  314 
nitride,  316 
oxide,  314 
phosphates,  315 
sulphate,  315 
Magnetite,  399 
Malachite,  335 
Manganese,  386 
alums,  384 
bronze,  336 
heptoxide,  39 1 
peroxide,  389 
Manganic  acid,  390 
compounds,  388 
Manganin,  405 
Manganites,  88,  389,  390 
Manganous  compounds,  387 
sulphate,  388 
Marsh  gas,  1 52 
test,  145 
Mass  action,  271 
Massicot,  369 
Matter,  18 

Maximum  valence,  170 
Mechanical  equivalent  of  heat,  22 
mixture,  18 

Mercuric  compounds,  325,  326 
Mercurous  compounds,  323,  324,  325 
Mercury,  321,  323,  323 
Meta-acids,  201 

j)hos])lioric  acid,  217 
stannic  acid,  367 
Metallic  ions,  402 
Metalloids  (non-metals),  19,  39 
oxygen  compounds  of,  172 
Mcrtals,  19,  39,  251 


Metals,  heavy,  252 
light,  252 

properties  of,  251,  255 
rare,  354 
Metargon,  116 
Methane,  152,  153 
Mineral  water,  91 
Mohr’s  salt,  398 
Mol,  98 

Mol-volume,  98 
Molecular  compounds,  170 
Molecules,  24,  25,  75,  167 
Molybdenum,  382 

compounds,  382,  383 
Molybdic  acid,  383 
Mosaic  gold,  367 
Mosandrium,  26 
Multiple  proportions,  71 
law  of,  72 
Muriatic  acid,  59 


N. 

Neodymium,  355 
Neon,  1 16 
Nickel,  404 
alloys,  405 
carbonyl,  233 
oxides,  405 
plating,  405 
sulphate,  405 
tetracarbonyl,  405 
Niobium,  226 
Nitramide,  204 
Nitrates,  202 
Nitric  acid,  202,  203 
fuming,  203 
oxide,  209,  210 
Nitrites,  202,  205 
Nitrogen,  115,  I16 
carbonyl,  234 
chloride,  133 
group,  1 14 

oxygen  derivatives  of,  201 
iodides,  134 
pentasulphide,  213 
pentoxide,  204 
tetroxide,  206 
trioxide,  205 
sulphide,  213 
Nitro-metals,  207 
Nitroso-acid,  208 
Nitrous  acid,  205 
oxide,  21 1 

Nitrosulphonic  acid,  207 
Nitrosyl  chloride,  204 
sulphuric  acid,  207 
Nitroxyl  chloride,  204 


INDEX. 


427 


Non-metals,  19,  39 
Nordhausen  sulphuric  acid,  194 
Normal  acids,  201 


O. 

Olefiant  gas,  154 
Organic  chemistry,  15 1 
Ortho-acids,  201 
Osmium,  410,  41 1 
Osmotic  pressure,  267 
Oxidation,  83 

theory  of  Traube,  loi 
Oxides,  84,  258 
indifferent,  84 
Oxygen,  80 

atomic  weight  of,  98 
group,  80 
liquid,  82 
oxidized,  87 
preparation  of,  80 
properties  of,  82 
reduction  of,  loi 
salts,  61 
Ozone,  84 

constitution  of,  86 
preparation  of,  85 
tests  for,  86 

thermo-chemical  deportment  of,  87 


P. 

Palladium,  412 

hydride,  46,  412 
Palladic  compounds,  412,  413 
Palladious  compounds,  412,  413 
Passive  iron,  397 
Pattinson’s  method,  337 
Pearl  ash,  281 
Pentathionic  acid,  199 
Perbromic  acid,  179 
Percarbonic  acid,  231 
Perchloric  acid,  178 

hydrate  of,  179 
Perchromic  acid,  374,  ^75 
anhydride,  381 
Periodates,  279 
Periodic  acid,  181 

hydrate  of,  181 
•sy.stem  of,  164,  243 
Periodicity  of  chemical  valence,  248 
Permanent  white,  309 
Permanganic  acid,  390 
Peroxides,  259 
Persulphates,  188 
Persulphuric  acid,  188 
Philippium,  26 
Phlogiston,  1 18 
Phosgene  gas,  234 


Phospham,  220 
Phosphates,  214 
Phosphine,  138,  139 

molecular  formula  of,  140 
Phosphites,  21 3 
Phosphonium,  140 
Phosphonium  iodide,  140 
Phosphoric  acid,  216 

anhydride,  217 
hypo-,  216 
meta-,  217 
pyro-,  216 
Phosphorite,  301 
Pho.sphorous  acid,  215 
oxide,  215 

Phosphorus,  135,  137 
acids,  214 
bromides,  142 
bronze,  336 
burns,  137 
chloranhydrides,  ,218 
fluoride,  142 
iodides,  142 
metallic,  136 
oxides,  213 
oxychloride,  218 
pentachloride,  141 
pentoxide,  217 
red,  136 
salt  of,  298 
sulphochloride,  219 
sulphur  derivatives,  220 
trichloride,  141 
yellow,  136 
Photography,  340 
Physical  phenomena,  18 
Pink  salt,  366 
Platinum,  413,  414 

ammonium-compounds,  415 
black,  413 

cyanogen  compounds  of,  415 
metals,  408,  409,  410 
heavy,  408 
light,  408 
separation  of,  409 
sponge,  413 

Platinic  compounds,  414 
Platinous  compounds.  414 
Plumbates,  369,  370 
Iflumbic  acid,  370 
Polychromates,  378 
Polysulphides,  283 
Polythionic  acids,  197 
Porcelain,  354 
Potashes,  281 
Potassium,  273 
borates,  280 
bromate,  278 
bromide,  276 


INDEX. 


428 


Potassium  carbonate,  280 
chlorate,  276 
cliloricle,  276 
chromates,  379 

compounds,  recognition  of,  283 
cyanide,  276 
ferrocyanide,  402 
ferricyanide,  403 
fluoride,  276 
hydride,  274 
hydroxide,  275 
hypochlorite,  278 
iodate,  278 
iodide,  276 
manganate,  390 
nitrate,  279 
nitrite,  280 
oxides,  275 
percarbonate,  282 
perchlorate,  278 
permanganate,  391 
phosphates,  280 
silicate,  282 
sulphates,  279 
sulphides,  282,  283 
sulphites,  279 
Praseodymium,  355 
Preparing  salts,  367 
Prussian  blue,  403 
Prussic  acid,  236 
Purple  of  Cassius,  345 
Pyrites,  400 
Pyroarsenic  acid,  222 
Pyromorphite,  367 
Pyrophosphoric  acid,  216 
Pyrosulphuric  acid,  193 
Pyrosulphuryl  chloride,  196 


Q- 

Quadrant  oxides,  259 
Quartz,  236 
(Quicksilver,  322 

R. 

Radiant  heat,  21 
Radical,  167 
Realgar,  223 
Reduction,  84 
Residue,  acid,  174 
Reversible  reaction,  91 
Rhodium,  4II 
Rimnann’s  green,  407 
Roentgen  rays,  21,  415 
Rose’s  metal,  373 
Rubidium,  284 
Ruby,  350 
Ruthenium,  410 


S. 

Safety  lamp,  159 
Salt  of  phosphorus,  298 
producers,  49 
Saltpeter,  279,  296 
Salts,  60,  259 

ammonium,  129 
basic,  260 
double,  260 
haloid,  61 
oxygen,  61 
transposition  of,  270 
Samarium,  356 
Sapphire,  350 
Scandium,  355 
Scheele’s  green,  335 
Schlippe’s  salt,  225 
Schweizer’s  reagent,  334 
Selenium,  1 12 
acids,  200 
chlorides,  112 
hydride,  1 1 2 
oxide,  199 
Siderite,  399 

Siemens-Martin  method,  397 
Silica,  236 
Silicates,  238 
Silicic  acid,  237,  238 
Silicon,  160 

bromide,  162 
bronze,  336 
carbide,  163 
chloride,  i6i 
chloroform,  162 
dioxide,  236 
disulphide,  238 

fluoride,  162  — _ 

hydride,  i6i 
iodide,  162 
Silver,  336,  337,  338 

allotropic  forms  of,  338 
bromide,  340 
chloride,  339 
coins,  331 
cyanide,  342 
fineness  of,  338 
iodide,  340 
nitrate,  34I 
nitride,  341 
nitrite,  341 
oxides,  338,  339 
plating,  342 
sulphate,  34I 
sulphide,  341 
Silvering,  342 
Slags,  396,  397 
Slaked  lime,  301 
Smalt,  406 
Soda,  caustic,  286 


INDEX. 


429 


Soda,  residue,  290 
Solvay,  291 
Sodium,  284 

azoimide,  295 
borate,  294 
bromide,  287 

carbonate,  289,  290,  291,  292 
chlorate,  287 
chloride,  286 

compounds,  recognition  of,  295 
Sodium  hydroxide,  286 
hyposulphite,  289 
iodate,  287 
iodide,  287 
nitrate,  293 
nitride,  295 
nitrite,  293 
oxides,  285 
perchlorate,  287 
periodate,  287 
peroxide,  285 
phosphates,  293,  294 
silicate,  295 
sulphate,  287,  288 
sulphite,  288 
thiosulphate,  289 
Soft  solder,  368 
Solutions,  265 

supersaturated,  288 
theory  of  dilute,  267 
thermo-chemistry,  92 
Special  sciences,  17 
Specific  gravity  of  gases,  73 
heat,  253 
volume,  45,  252 

Spectrum  analysis,  415,  416,  417,  418, 
419 

Speiss-cobalt,  406 
Spinels,  351 
Stannic  acid,  66,  367 

compounds,  366,  367 
Stannous  compounds,  365 
Status  nascens,  53,  77,  88,  102,  130,  175, 
203 

Steel,  394 
Stibine,  147 
Stoechiometric  laws,  71 
S trass,  306 
Strontium,  307 

carbonate,  308 
nitrate,  308 
oxide,  307 

Structure,  chemical,  168 
Suboxides,  259 
Substitution,  159 
Sulphamide,  196 
Sulphates,  193 
Sulphides,  108 
Sulphimide,  196 


Sulphites,  186 
Sulphocarbonic  acid,  235 
Sulpho-group,  198 
Sulpho-stannates,  367 
Sulphur,  104,  105,  106 
bromides,  iii 
chemical  properties  of,  106 
chlorides,  no,  iii 
dioxide,  183,  184 
heptoxide,  188 
iodides,  iii 

oxygen  compounds  of,  183 
sesquioxide,  187 
trioxide,  187 

Sulphuric  acid,  189,  190,  19 1,  192,  193 
amides,  196 
chloranhydrides,  195 
Sulphuric  acid,  di-,  193 
fuming,  194 
hydrates,  192 
Nordhausen,  194 
Sulphurous  acid,  185 
Sulphuryl,  195 
chloride,  185 
Superphosphate,  305 
Sylvite,  276 
Sympathetic  ink,  406 


T. 

Tantalum,  226 
Tellurium,  113 
acids,  200 
bromides,  113 
chlorides,  113 
hydride,  1 13 
Tension  of  vapors,  90 
Terbium,  26,  356 
Tetrathionic  acid,  199 
Thallic  acid,  360 

compounds,  360 
Thallium,  359 
alum,  360 

Thallous  compounds,  360 
Thenard’s  blue,  407 

Thermo-chemistry  of  the  elements,  28,  66, 
92,  103,  109,  1 14,  129,  133,  134,  137, 
142,  160,  164,  186,  187,  212,  236,  327, 
328,  330,  340,  348,  387 
Thionic  acids,  I97 
Thionyl  chloride,  185 
Thiosulphuric  acid,  197 
Thomas  slag,  397 
Thorium,  238,  240 
Thulium,  27,  356 
Thyroidin,  55 
Tin,  164,  238,  363 
dichloride,  364 
dioxide,  366 


INDEX. 


430 

'l  in  (lisulpliide,  367 
monoxide,  365 
oxysulphide,  367 
salt,  364 
stone,  363 
tetrachloride,  366 
Titanium,  238,  239 
Trithionic  acid,  199 
d'ornbac,  336 
Tungsten,  384 

compounds,  384 
Turpeth  mineral,  327 
Turnbull’s  blue,  403 
Type  metal,  368 


U. 

Uchatius  steel,  397 
Ultramarine,  354 
Uranium,  384 
Uranic  compounds,  385 
Uranous  compounds,  385 
Uranyl,  385 


V. 

Valence,  165,  166,  167,  169 
periodicity  of,  248 
variable,  170 
Vanadium,  226 
Varec,  54 

Vitriol,  copper,  334 
green,  398 

oil  of  (see  Sulphuric  acid) 
Volume,  atomic,  252 
specific,  252 


W. 

Water,  88 

chemical  properties  of,  9 1 
crystallization,  266 
constitution.  315 
dissociation  of,  93 
distilled,  91 


Water,  electrolysis  of,  92 
gas,  232 
glass,  295 
hard,  91 
mineral,  91 

molecular  formula  of,  97 
natural,  91 
oxidation  of,  lOO 
physical  i)roperties  of,  88 
({uantitative  composition  of,  95 
soft,  91 

thermo-chemical  deportment  of,  92 
Weight  proportions  in  the  union  of  the  ele- 
ments, 69 

Welsbach  mantle,  240 
White  precipitate,  325 
lead,  371 

Wood’s  metal.  373 
Wulfenite,  3^7 


X. 

Xenon,  1 16 

V. 

Yellow  prussiate  of  potash,  402 
Ytterbium,  356 
Yttrium,  355 


Z. 

Zinc,  318 

blende,  319 
chloride,  319 
dust,  318 
oxide,  318 
silicate,  319 
sulphate,  319 
sulphide,  319 
w'hite,  318 
Zircon,  239 
Zirconium,  238,  239 
light,  83 


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INORGANIC  CHEMISTRY.  Fifth  American  from  the 
Tenth  German  Edition  by  Prof.  H.  Klinger,  University  of 
Konigsberg.  Thoroughly  Revised  and  in  many  parts  Rewritten. 
With  many  Illustrations  and  Colored  Plate  of  Spectra.  i2mo. 

Cloth,  gi.75 

I'rom  The  Scientific  American. 

” It  is  markedly  practical,  clear,  and  direct  in  its  statements,  bringing  out  prominently 
the  relations  between  proved  facts  and  theories  or  hypotheses,  so  as  to  preclude  as  far  as 
))OSsible  si)eculative  inferences  from  the  mind  of  a student  beginning  the  study  of 
chemistry.” 


4 


BOOKS  FOR  CHEMISTS  AND  MANUFACTURERS. 

CONGDON.  Laboratory  Instructions  in  General  Chem- 
istry. By  Ernest  A.  Congdon,  Professor  of  Chemistry  in  the 
Drexel  Institute,  Philadelphia;  Member  American  Chemical 
Society ; Fellow  London  Chemical  Society,  etc.  Illustrated. 

In  Press. 

ULZER  AND  FRAENKEL.  Introduction  to  Chemical- 
Technical  Analysis.  By  Prof.  F.  ULZERand  Dr.  A.  Fraenkel, 
Directors  of  the  Testing  Laboratory  of  the  Royal  Technological 
Museum,  Vienna.  Authorized  IVanslation  by  Hermann  Fleck, 
NAT.sc.D.,  Instructor  in  Chemistry  in  the  John  Harrison  Laboratory 
of  Chemistry,  University  of  Pennsylvania.  With  an  A]}i)endix  by 
the  Translator  relating  to  I'ood-Stuffs,  Asphaltum,  and  Paint.  12 
Illustrations.  8vo.  Cloth,  $1.25 

With  a view  to  increasing  its  usefulness  as  a text-book  and  laboratory  guide, 
the  translator  has  added  descriptions  of  the  analysis  of  typical  food  stuffs  and  other 
important  products. 

BLOXAM.  Chemistry — Inorganic  and  Organic.  With 
Experiments.  By  Charles  L.  Bloxam.  Edited  and  Revised  by 
J.  M.  Thompson,  Professor  of  Chemistry  in  King’s  College,  Lon- 
don, and  A.  G.  Bloxam,  Head  of  Chemistry  Department,  The 
Goldsmith  Institute,  London.  Eighth  Edition,  Enlarged.  281  En- 
gravings, 20  of  which  are  new.  8vo.  Cloth,  ^4. 25;  Leather,  $5. 25 

This — the  eighth — edition  has  been  very  thoroughly  revised.  Many  changes 
have  been  made,  though  the  general  arrangement  remains  the  same.  The  sec- 
tions on  explosives,  to  which  the  work  owes  considerable  reputation,  have  been 
carefully  gone  over.  It  is  in  some  respects  more  thorough  and  complete  than  any 
other  work. 

BARTLEY.  Medical  and  Pharmaceutical  Chemistry. 

By  E.  H.  Bartley,  m.d..  Professor  of  Chemistry  and  Toxicology 
at  the  Long  Island  College  Hospital ; Dean  and  Professor  of 
Chemistry,  Brooklyn  College  of  Pharmacy;  Chief  Chemist,  Board 
of  Health  of  Brooklyn.  Fifth  Edition,  Revised  at.d  Enlarged. 
92  Illustrations  and  Glossary.  121110.  Cloth,  $3.00 

“This  excellent  book  contains,  in  the  most  concise  form,  all  the  knowledge 
of  medical  and  pharmaceutical  chemistry.  The  present  edition  has  been  greatly 
enlarged,  and  a new  chapter  on  Physiological  and  Clinical  Chemistry  added.  . . . 
The  whole  book  reads  admirably  well  and  deserves  the  highest  recommenda- 
tion.”— N.  Y.  Medical  Record. 

Clinical  Chemistry.  The  Chemical  Examination  of  the 
Saliva,  Gastric  Juice,  Feces,  Milk,  Urine,  etc.,  with  notes  on 
Urinary  Diagnosis,  Volumetric  Analysis,  and  Weights  and  Meas- 
ures. Illustrated.  121110.  Cloth,  $1.00 


The  prices  of  these  books  are  net.  No  discount  allowed  retail  purchasers. 


BOOKS  FOR  CHEMISTS  AND  MANUFACTURERS. 


0 


TRAUBE.  Physico-Chemical  Methods.  By  Dr.  J.  Traube, 
Privatdocent  in  the  Technical  High  School  of  Berlin.  Authorized 
Translation  by  W.  L.  Hardin,  Ph.D.,  Instructor  in  Chemistry, 
University  of  Pennsylvania.  97  Illustrations.  8vo.  Cloth,  $1.50 

“ Dr.  Traube’s  book  should  certainly  be  welcomed  by  us  all  on  account  of 
the  clear  and  concise  exposition  he  has  given  us  of  the  most  important  of  physico- 
chemical methods.  The  value  of  this  book  is  further  enhanced  by  a number  of 
good  cuts  which  aid  one  in  obtaining  a clear  idea  of  the  methods  under  discus- 
sion. The  book  has  been  carefully  edited,  and  the  typography  is  good.” — The 
Journal  of  the  Worcester  Polytechnic  Institute, 

CALDWELL.  Chemical  Analysis.  Elements  of  Quali- 
tative and  Quantitative  Chemical  Analysis.  By  G.  C.  Caldwell, 
B.s. , PH.D.,  Professor  of  Agricultural  and  Analytical  Chemistry  in 
Cornell  University,  Ithaca,  New  York,  etc.  Third  Edition.  Re- 
vised and  Enlarged.  Octavo.  Cloth,  ^1.50 

This  book  has  been  extensively  used  in  American  Schools  and  Colleges,  and 
needs  no  special  introduction.  Written  to  supply  Prof.  Caldwell’s  need  in  his 
own  work,  it  is  an  eminently  practical  and  complete  text-book. 

HOLLAND.  The  Urine,  the  Gastric  Contents,  the 
Common  Poisons,  and  the  Milk.  Memoranda,  Chemical  and 
Microscopical,  for  Laboratory  Use.  By  J.  W.  Holland,  m.d.. 
Professor  of  Medical  Chemistry  and  Toxicology  in  Jefferson  Med- 
ical College,  Philadelphia.  Sixth  Edition,  Enlarged.  Illustrated 
and  Interleaved.  i2mo.  Cloth,  ^i.oo 

LEFFMANN.  Compend  of  Medical  Chemistry,  Inor- 
ganic and  Organic.  Including  Urine  Analysis,  Chemistry  of  Tissues 
and  Secretions,  Table  of  Symbols,  Valencies,  Atomic  Weights,  etc. 
By  Henry  Leffmann,  m.d..  Professor  of  Chemistry  and  Metal- 
lurgy in  the  Pennsylvania  College  of  Dental  Surgery  and  in  the 
Wagner  Free  Institute  of  Science,  Philada.;  Professor  of  Chemistry, 
Woman’s  Medical  College  of  Pennsylvania,  etc.  Fourth  Edition. 

Cloth,  .80;  Interleaved  for  the  addition  of  Notes,  $1.25 

The  Coal-Tar  Colors,  with  Special  Reference  to  their 
Injurious  Qualities  and  the  Restrictions  of  their  Use.  A Trans- 
lation of  Theodore  Weyl’s  Monograph.  121110.  Cloth,  ^1.25 

Structural  Formulae  for  the  Use  of  Students.  Contain- 
ing 180  Structural  and  Stereo-Chemic  Formulae.  121110.  Inter- 
leaved. Cloth,  ^i.oo 

***See  also  Water  and  Milk  Analyses,  page  lo. 


The  prices  of  these  books  are  net.  No  discount  allowed  retail  purchasers. 


(y  BOOKS  FOR  CHEMISTS  AND  MANUFACTURERS. 


CLOWES  AND  COLEMAN.  Quantitative  Analysis. 

Adapted  for  the  Use  of  the  Laboratories  of  Schools  and  Colleges. 
By  Frank  Clowes,  sc.d.,  Emeritus  Professor  of  Chemistry,  Univer- 
sity College,  Nottingham,  and  I.  Bernard  Coleman,  Assoc.  R.  C. 
Sci.,  Dublin,  Professor  of  Chemistry,  Southwest  London  Polytech- 
nic. Fifth  Edition.  1 1 7 Illustrations.  Neaj'ly  Ready. 

OETTEL.  Practical  Exercises  in  Electro-Chemistry. 

By  Dr.  Felix  Oettel.  Authorized  Translation  by  Edgar  F. 
Smith,  m.a..  Professor  of  Chemistry,  University  of  Pennsylvania; 
Translator  of  Richter’s  Chemistries,  etc.  Illustrated.  Cloth,  .75 
Introduction  to  Electro-Chemical  Experiments.  Illus- 
trated. By  same  Author  and  Translator.  Cloth,  .75 

SMITH.  Electro-Chemical  Analysis.  By  Edgar  F.  Smith, 
Professor  of  Chemistry,  University  of  Pennsylvania.  Second 
Edition.  Enlarged.  28  Illustrations.  i2mo.  Cloth,  ^1.25 

***  This  book  has  recently  been  translated  into  German  by  Dr.  Max  Ebeling 
of  Berlin,  and  into  French  by  Prof.  J.  Rosset. 

“ The  greatest  advantage  of  the  electrolytic  method  of  analysis  is  its  sim- 
plicity. We  do  not  require  to  introduce  materials  the  purity  of  which  often  leaves 
something  to  be  desired,  and  which  have  to  be  again  removed,  frequently  at  the 
cost  of  much  time  and  trouble.  Hence,  where  available,  the  electrolytic  method 
of  analysis  possesses  an  indisputable  advantage.  We  hope  that  Prof.  Smith’s 
little  work  will  call  increased  attention  to  this  branch  of  mineral  analysis.”-— 
Chemical  News,  London. 

' “ Chemists  will  find  this  little  book  an  excellent  guide  to  a knowledge  of  the 
methods  of  quantitative  analysis  by  electrolysis.” — American  Chemical  Journal, 
Baltimore,  Md. 

MUTER.  Practical  and  Analytical  Chemistry.  By 

John  Muter,  f.r.s.,  f.c.s.,  etc.  Second  American  from  the 
Eighth  English  Edition.  Revised  to  meet  the  requirements  of 
American  Medical,  Pharmaceutical,  and  Dental  Colleges.  56  Illus- 
trations. Cloth,  ^1.25 

“ The  editor  has  made  only  such  changes  as  were  required  to  adapt  the  book 
to  the  U.  S.  Pharmacopoeia,  except  in  the  chapter  on  urine  analysis,  which  has 
been  enlarged  and  to  which  cuts  of  microscopic  sediments  and  other  illustrations 
have  been  added.  Several  other  processes  have  been  added,  such  as  estimation 
of  chloral  hydrate,  of  fat  in  milk,  etc.” — The  Popular  Science  Monthly. 

TYSON.  Guide  to  the  Practical  Examination  of  Urine. 

Ninth  Edition.  By  James  Tyson,  m.d..  Professor  of  Medicine, 
University  of  Pennsylvania,  etc.  With  Colored  Plate  and  48 
Illustrations.  Ninth  Edition.  12010.  276  pages.  Cloth,  ;^i. 25 


The  prices  of  these  books  are  net.  No  discount  allowed  retail  purchasers. 


BOOKS  FOR  CHEMISTS  AND  MANUFACTURERS. 


7 


SMITH  AND  KELLER.  Experiments.  Arranged  for 
Students  in  General  Chemistry.  By  Edgar  F.  Smith,  Professor  of 
Chemistry,  University  of  Pennsylvania,  Translator  of  Richter’s 
Chemistries,  etc.,  and  Dr.  H.  F.  Keller,  Professor  of  Chemistry, 
Philadelphia  High  School.  Third  Edition.  41  Ulus.  Cloth,  .60 

“This  series  of  exercises,  based  on  the  authors’  experiences  with  their  own 
classes,  is  intended  to  accompany  any  convenient  text-book  of  inorganic  chemistry ; 
but  reference  is  made  to  that  of  Richter,  Beginning  with  fundamental  operations 
(as  with  blowpipe,  glass  tubing,  balance,  and  graduates)  and  general  principles 
(as  the  difference  between  chemical  and  physical  change),  the  course  proceeds  to 
the  study  of  hydrogen  and  other  non-metals  in  Part  I,  followed  by  the  metals  in 
Part  II.  Quantitative  relations  are  well  presented  in  the  experimental  work  and 
stoichiometrical  problems,  as  in  determining  the  ^/-equivalent  of  zinc,  the  density 
of  Cl,  eudiometric  combustion  of  methane,  etc.” — Science. 

THRESH.  Water  and  Water  Supplies.  By  John  C. 
Thresh,  d.sc.  (Eondon),  m.d.,  d.ph.  (Cambridge),  Medical 
Officer  of  Health  to  the  Essex  County  Council ; Lecturer  on  Public 
Health,  King’s  College,  London ; Fellow  of  the  Institute  of 
Chemistry,  etc.  Second  Edition,  Revised.  Illustrated.  i2mo. 
438  pages.  Cloth,  ^2.00 

SUTTON.  Volumetric  Analysis,  a Systematic  Handbook 
of ; or,  the  Quantitative  Estimation  of  Chemical  Substances  by 
Measure,  applied  to  Liquids,  Solids,  and  Gases.  Adapted  to  the 
requirements  of  Pure  Chemical  Research,  Pathological  Chemistry, 
Pharmacy,  Metallurgy,  Manufacturing  Chemistry,  Photography, 
etc.,  and  for  the  Valuation  of  Substances  used  in  Commerce,  Agri- 
culture, and  the  Arts.  By  Francis  Sutton,  f.c.s.,  f.i.c..  Public 
Analyst  for  the  County  of  Norfolk.  Eighth  Edition.  Revised 
and  Improved.  112  Illustrations.  Just  Ready.  Cloth,  $5.00 

“ We  need  only  say  that  this  work  is  welcome  as  bringing  ‘ Sutton’s  Volu- 
metric Analysis  ’ down  to  date.  The  original  book  has  been  so  long  the  standard, 
and  is  so  well  known,  that  description  seems  hardly  necessary.” — The  Scientific 
American. 

SYMONDS.  Manual  of  Chemistry,  for  Medical  Students. 
By  Brandreth  Symonds,  a.m  , m.d.,  Ass’t  Physician  Roosevelt 
Hospital,  Out-Patient  Department;  Attending  Physician  North- 
western Dispensary,  New  York.  Second  Edition.  Cloth,  $2.00 

SMITH.  Dental  Metallurgy.  A Manual.  By  Ernest  A. 
Smith,  f c s..  Assistant  Instructor  in  Metallurgy,  Royal  College 
of  Science,  London.  Illustrated.  i2mo.  Cloth,  $1.75 

“The  metals  in  common  use  in  dental  practice  are  each  described,  and  a 
special  chapter  is  devoted  to  the  question  of  Dental  Amalgams.  With  a view  of 
rendering  the  book  as  practical  as  possible,  the  author  has  enumerated  a list  of 
100  experiments.” — llie  Lancet,  London. 


The  prices  of  these  books  are  net.  No  discount  allowed  retail  purchasers. 


s BOOKS  FOR  CHEMISTS  AND  MANUFACTURERS. 


WOODY.  Essentials  of  Medical  and  Clinical  Chemistry. 
With  Laboratory  Exercises,  by  Sam  E.  Woody,  a.m.,  m.d.,  I’ro- 
fessor  of  Clieniistry  and  Diseases  of  Cliildren  in  tlie  Medical 
De])artment  of  Kentucky  University,  Louisville.  Fourth  lOdition. 
Revised,  with  Tables  and  Illustrations.  121110.  Nearly  Ready. 


TOXICOLOGY  AND  MED.  JURISPRUDENCE. 

REESE’S  Medical  Jurisprudence  and  Toxicology.  A 

Text-book  for  Medical  and  Legal  Practitioners  and  Students.  By 
John  J.  Reese,  m.d.,  late  Professor  of  the  Priiiciides  and  Practice 
of  Medical  Jurisprudence,  including  Toxicology,  in  the  University 
of  Pennsylvania,  Medical  Department.  Fifth  Edition.  Revised 
and  Enlarged  by  Henry  Leffmann,  m.d..  Pathological  Chemist, 
Jefferson  Medical  College  Hospital,  Philadelphia;  Hygienist  and 
Food  Inspector  Pennsylvania  State  Board  of  Agriculture,  etc. 
121110.  646  jiages.  Cloth,  ^3.00;  Leather,  ^3.50 

“ To  the  student  of  medical  jurisprudence  and  toxicology  it  is  invaluable,  as 
it  is  concise,  clear,  and  thorough  in  every  respect.” — The  American  Journal  of 
the  Medical  Sciences. 

“ The  book  happily  meets  the  needs  of  students,  and  we  unqualifiedly  com- 
mend it.” — American  Practitioner  and  News,  Louisville. 

“ The  book  will  be  found  to  be  a useful  one,  and  as  such  we  commend  it  to 
students  of  law  and  medicine.” — Marshall  D.  Ewell,  Dean  of  the  Kent  Law 
School,  Chicago. 

MANN.  Forensic  Medicine  and  Toxicology.  A Text- 
book by  J.  Dixon  Mann,  m.d.,  f.r.c.p..  Professor  of  Medical 
Jurisprudence  and  Toxicology  in  Owens  College,  Manchester; 
Examiner  in  Forensic  Medicine  in  University  of  London,  etc. 
Illustrated.  Octavo.  Cloth,  ^6.50 

TANNER’S  Memoranda  of  Poisons  and  their  Antidotes 
and  Tests.  By  Thos.  Hawkes  Tanner,  m.d.,  f.r.c.p.  Seventh 
American,  from  the  last  London  Edition.  Revised  by  John  J. 
Reese,  m.d..  Professor  of  Medical  Jurisprudence  and  Toxicology 
in  the  University  of  Pennsylvania.  i2mo.  Cloth,  .75 

“The  fact  of  any  technical  work,  great  or  small,  reaching  Its  seventh  edition 
speaks  for  itself.  In  the  most  condensed  form  are  given  the  history  of  poisons,  their 
antidotes,  and  various  mechanical  methods  for  overcoming  the  tendency  toward 
death.  'I'he  principal  cnanges  in  the  new  edition  have  been  the  substitution  of 
modern  chemical  nomenclature  and  the  omission  of  obsolete  portions  of  the  old 
text.  I'he  toxicology  of  poisonous  food  has  been  presented  as  fully  as  the  concise 
character  of  the  book  allows.'’ — Medical  Record,  A^e^v  York. 


The  prices  of  these  books  are  net.  No  discount  allowed  retail  purchasers. 


TECHNOLOGICAL  BOOKS. 


9 


GROVES  AND  THORP.  Chemical  Technology.  The 

Application  of  Chemistry  to  the  Arts  and  Manufactures.  Edited 
by  Charles  E.  Groves,  f.r.s.,  and  Wm.  Thorp,  b.Sc.,  f.i.c. 
With  numerous  Illustrations.  Each  volume  sold  separately. 

Vol.  I.  Fuel  and  Its  Applications.  By  Dr.  E.  J.  Mills, 
F.R.S.,  Professor  of  Chemistry,  Anderson  College,  Glasgow;  and 
Mr.  F.  J.  Rowan,  c.e.,  assisted  by  an  American  expert.  607 
Illustrations  and  4 Plates.  Cloth,  ^5.00;  Half  Morocco,  $6.50 

“ It  is  without  doubt  the  most  useful  and  comprehensive  book  in  the  English 
language  on  fuels,  and  is  a valuable  acquisition  to  our  standard  books  of  refer- 
ence.”— Journal  of  the  Frajiklin  Institute. 

“ It  covers  a wide  range  of  knowledge,  and  should  be  at  the  elbow  of  every 
intelligent  and  progressive  manufacturer.  ” — The  Iron  Trade  Review,  Cleveland. 

“ The  book  will  be  very  useful  for  reference,  and  should  be  of  especial  value 
to  the  inventors  and  experimenters  or  users  of  processes  or  appliances  for  the  com- 
bustion of  fuels,  since  in  it  can  be  found  a record  of  a large  part  of  the  methods 
heretofore  proposed  and  adopted.  Where  critical  remarks  are  made  they  appear  to 
be  judicious.  The  illustrations  are  very  numerous  and  are  well  selected.  An 
immense  amount  of  information  has  been  crowded  into  these  closely  printed  802 
pages.” — Engineering  and  Alining  Journal,  Afeiv  York. 

“ The  book  is  very  fully  illustrated,  as,  indeed,  the  nature  of  the  subject 
requires,  and  includes  a large  number  of  tables  giving  fuel  statistics,  analyses  of 
different  fuels,  and  comparative  results.  ” — The  Railroad  and  Engineering  Journal. 

Vol.  II.  Lighting.  Candles,  Oils,  Lamps,  etc.  By  W.  Y. 
Dent,  I.  McArthur,  L.  Field,  F.  A.  Field,  Boverton  Red- 
wood, and  D.  A.  Louis.  358  Illustrations.  Octavo. 

Cloth,  $4.00;  Half  Morocco,  ^5.50 
Vol.  III.  Gas  Lighting.  By  Charles  Hunt,  Manager  of 
the  Birmingham  Corporation  Gasworks.  Illustrated.  Octavo. 

Cloth,  $3.50  ; Half  Morocco,  ^4.50 
Vol.  IV.  Electric  Lighting.  Pt^eparhig, 

GARDNER.  The  Brewer,  Distiller,  and  Wine  Manu- 
facturer. Giving  full  Directions  for  the  Manufacture  of  Beers, 
Spirits,  Wines,  Liquors,  etc.,  etc.  A Handbook  for  all  interested 
in  the  manufacture  and  sale  of  Alcohol  and  Its  Compounds.  Edited 
by  John  Gardner,  f.c.s..  Editor  of  '‘Cooley’s  Cyclopedia”  and 
“ Beasley’s  Druggists’  Receipt  Book.”  Illustrated.  Cloth,  ^1.50 

“Trustworthy  and  valuable.” — German  and  America}!  Brezvers'  Journal. 

“A  very  complete  handbook.” — Boston  Journal  of  Chemistry. 

Bleaching,  Dyeing,  and  Calico  Printing.  With  P'ormulae ; 
a Chapter  on  Dye  Stuffs.  Illustrated.  121110.  Cloth,  $1.50 

“A  serviceable  manual.” — Inventors'  and  Manufacture}^  Gazette. 


The  prices  of  these  books  are  net.  No  discount  allowed  retail  purchasers. 


10  BOOKS  FOR  CHEMISTS  AND  MANUFACTURERS. 


CAMERON.  Oils  and  Varnishes.  A Practical  Handbook, 
by^  James  Cameron,  f.i.c.  With  Illustrations,  Formiiloe,  d’ables, 
etc.  i2mo.  Cloth,  $2.25 

Soap  and  Candles.  A New  Handbook  for  Manufacturers, 
Chemists,  Analysts,  etc.  54  Illustrations.  121110.  Cloth,  $2.00 


EMERGENCIES. 

DULLES.  Accidents  and  Emergencies.  A Manual  for 
the  Treatment  of  Surgical  and  other  Injuries,  Poisoning  and  various 
Domestic  Emergencies,  in  the  absence  of  the  Ph}'sician.  By 
Charles  W.  Dulles,  m.d.,  Surgeon  to  the  Out-Door  Department 
of  the  University  and  Presbyterian  Hospitals,  Philadelphia.  Fifth 
Edition,  Enlarged.  New  Illustrations.  121110.  Cloth,  ^i.oo 

“This  is  a revised  and  enlarged  edition,  with  new  illustrations,  of  the 
manual,  explaining  the  treatment  of  surgical  and  other  injuries  in  the  absence  of 
the  physician.  The  simple  and  practical  suggestions  of  this  little  book  should  be 
known  to  every  one.  Accidents  are  constantly  occurring,  and  a knowledge  of 
what  should  be  done  in  an  emergency  is  very  valuable.  Such  a handbook  should 
be  in  every  home,  placed  where  it  can  always  be  found  readily,” — Boston  Joiirnal 
of  Education. 


WATER  AND  MILK  ANALYSES. 

LEFFMANN.  Examination  of  Water  for  Sanitary  and 
Technical  Purposes.  Presenting  those  Processes  that  are  Most 
Trustworthy  and  Practical.  By  Henry  Leffmann,  m.d..  Professor 
of  Chemistry  and  Metallurgy,  Pennsylvania  College  of  Dental 
Surgery;  Professor  of  Chemistry,  Woman’s  Medical  College  of 
Pennsylvania,  etc.  Fourth  Edition.  Revised  and  Enlarged.  Illus- 
trated. i2mo.  Cloth,  $1.25 

“ This  is  a well-compiled  and  useful  little  treatise.” — London  Lancet. 

“ An  admirable  digest  of  our  present  knowledge.”— of  Analytical 
Chemistry. 

“ Especially  valuable  is  the  section  on  interpretation  of  results.” — Railroad 
and  Ejigineering  Neivs. 

BY  THE  SAME  AUTHOR. 

Analysis  of  Milk  and  Milk  Products.  Arranged  to  suit 
the  needs  of  Analytical  Chemists,  Dairymen,  and  Milk  Inspectors. 
Second  Edition.  Enlarged.  121110.  Cloth,  $1.25 

“ 'fhe  book  is  one  which  will  be  useful  in  the  hand  of  the  dairyman,  as  well 
as  in  the  hands  of  those  whose  duty  it  is  to  see  that  he  deals  fairly  with  his  cus- 
tomers.”— London  Sanitary  Record. 


The  prices  of  these  books  are  net.  No  discount  allowed  retail  purchasers. 


BOOKS  FOR  SCIENTISTS. 


11 


CARPENTER.  The  Microscope  and  Its  Revelations.  By 

W.  B.  Carpenter,  m d , f r s.  Eighth  Edition.  By  Rev.  Dr. 
Dai.ltnger,  f r s.  Revised  and  Enlarged,  with  800  Illustrations 
and  21  Lithographs.  Octavo.  1100  pages.  Pi'eparing. 

“ The  book,  therefore,  cannot  fail  to  be  of  value  to  chemists  and  others  at 
iron  works  intrusted  with  the  microscopical  examination  of  the  metals  with  which 
they  deal.” — The  Americaii  Manufacturer,  Pittsburgh. 

“ Every  one  who  has  a microscope  will  need  also  Carpenter’s  book  to  get  the 
most  out  of  his  instrument,  and  every  one  who  has  the  book  will  be  certain  to  want 
a microscope.” — Popular  Science  Monthly. 

“It  is  without  a rival  in  its  particular  field,  and  is  beyond  question  the  best 
single  work  on  the  subject,  not  only  in  English  but  in  any  other  language.  . . . 

A splendid  specimen  of  the  book-maker’s  art.” 

“ The  book  is  more  than  ever  a standard,  unrivaled  in  its  kind,  and  is  a neces- 
sity to  every  one  who  pretends  to  any  scientific  use  of  the  microscope.” — New  York 
Evening  Post. 

WETHERED.  Medical  Microscopy.  By  Frank  J.  Weth- 
ered,  M.D.,  M.R.c.p.  With  100  Illustrations.  121110. 

Cloth,  $2.00 

REEVES.  Medical  Microscopy.  Including  chapters  on 
Bacteriology,  Neoplasms,  Urinary  Examination,  etc.  By  James  E. 
Reeves,  m.d.,  Ex-President  American  Public  Health  Association, 
etc.  Illustrated.  121110.  Cloth,  $2.50 


THE  BEST  DICTIONARY. 

GOULD.  Illustrated  Dictionary  of  Medicine  and  Allied 
Sciences,  including  Chemistry,  Biology  (Zoology  and 
Botany;,  Hygiene,  etc.  Large,  Square  Octavo.  1633  pages. 

Fulfsheq;?,^^^^  } Thumb  Index,  $11.00 

Fifth  Edition.  Half  Russia,  Thumb  Index,  $12.00 

***  There  being  no  special  dictionaries  devoted  to  Chemistry  and  Biology, 
it  was  thought  eminently  proper  to  include  both  these  sciences  in  this  book.  They 
are  closely  related  to  medicine,  and  each  is  largely  dependent  upon  the  others, 
(jould’s  Illustrated  Ihctionary  contains  much  special  information  of  practical  use  to 
the  general  scientist.  Bacteriology  and  Parasitology  are  particularly  well  pre- 
sented, while  the  numerous  tables  of  Acids,  Alcohols,  Aldehyds,  Carbohydrates, 
Electric  Units,  Ethers,  P’oods,  Hydrocarbons,  Laws,  Milks,  Oils,  Pigments,  Pto- 
mains,  Resins,  Soaps,  Stains,  Starches,  Sugars,  Tests,  Theories,  Wave  Lengths, 
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12 


HYGIENE  AND  SANITARY  SCIENCE. 


HOTTER  and  FIRTH.  The  Theory  and  Practice  of 
Hygiene.  A Complete  Treatise  by  J.  Lane  Notter,  m.a.,  m d,, 
F.C.S.,  Fellow  and  Member  of  Council  of  the  Sanitary  Institute  of 
Great  Britain;  Jh'ofessor  of  Hygiene,  Army  Medical  School; 
Examiner  in  Hygiene,  University  of  Cambridge,  etc  , and  R.  H. 
Firth,  f.r.c.s.,  Assistant  Professor  of  Hygiene,  Army  Medical 
School,  Netly.  Illustrated  by  lo  Lithographic  Plates  and  other 
Illustrations,  and  including  many  Useful  Tables.  Second  Edition, 
Carefully  Revised.  Octavo.  In  Press. 

The  standard  authority  in  tlie  United  States  and  british  Army  and  Navy. 


» 

I 


PARKES.  Hygiene  and  Public  Health.  A Practical 
Manual.  By  Louis  C.  Parkes,  m.d.,  d.p.h.,  London  Hospital  ; 
Assistant  Professor  of  Hygiene  and  Public  Health  at  University 
College,  etc.  Fifth  Edition,  Enlarged  and  Revised!  8o  Illustra- 
tions. 121110.  Cloth,  $2.50 

“ Dr.  Parkes’  experience  as  a teacher  of  public  hygiene  has  enabled  him  to 
deal  with  the  whole  subject  in  a practical  and  intelligible  manner.” — The  Boston 
Medical  and  Surgical  Journal. 

STEVENSON  AND  MURPHY.  A Treatise  on  Hy- 
giene.  Illustrated.  Edited  by  Thomas  Stevenson,  m.d.,  f.r.c.p., 
Lecturer  on  Chemistry  and  on  Medical  Jurisprudence  at  Guy’s  Hos- 
pital, Official  Analyst  to  the  Home  Office ; and  Shirley  F.  Mur- 
phy, Medical  Officer  of  Health  to  the  County  of  London. 

Vol.  I.  9 Plates.  186  Illustrations.  1013  pages.  Cloth,  56.00 
Vol.  II.  45  Plates.  31  Illustrations.  847  pages.  Cloth,  56.00 
Vol.  III.  Sanitary  Law.  459  pages.  Cloth,  55.00 

EACH  VOLUME  SOLD  SEPARATELY. 

“ The  different  topics  are  fully  and  intelligently  treated,  especially  those 
•which  relate  to  the  subjects  of  Ventilation,  Water,  Soil,  Food,  Physical  Education, 
the  Dwelling,  and  the  Disposal  of  Refuse.  The  work  is  fully  illustrated  with 
plates,  diagrams,  and  wood-cuts,  and  pains  appear  to  have  been  taken  to  bring  the 
information  upon  each  topic  up  to  date.” — The  Boston  Medical  and  Surgical 
Journal. 

“ All  the  topics  are  treated  with  a thoroughness  of  detail  leaving  nothing  to 
be  desired.  The  contents  are  valuable  alike  to  the  physician,  the  municipal  health 
officer,  and  the  sanitary  engineer.” — Medical  Record.,  New  York. 

KENWOOD.  Public  Health  Laboratory  Work.  By  H. 

R.  Kenwood,  m.il,  d.p.h.,  f.c.s..  Instructor  in  Hygienic  Labora- 
tory, University  College;  late  Assistant  Examiner  in  Hygiene, 
Science  and  Art  Department,  South  Kensington,  London,  etc. 
With  1 16  Illustrations  and  3 Plates.  Cloth,  52.00 

***For  complete  list  of  books  on  Hygiene,  send  for  our  Special  Catalogue. 

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